•BiBHBBi 


GIFT  OF 
Agriculture   e'ducation 


ANTOINE  LAURENT  LAVOISIER  (1743-1794) 

Founder  of  modern  chemistry.    The  cut  is  from  a  picture  of  a  monument 
erected  in  his  honor  in  Paris.   Just  back  of  it  are  seen  the  pillars  of  the  church 
of  St.  Mary  Magdalene.   The  erection  of  this  church  was  begun  in  1806  by 
Napoleon  I,  in  commemoration  of  his  victories 


FIRST  COURSE  IN 
CHEMISTRY 


BY 
WILLIAM  McPHERSON 

AND 

WILLIAM  EDWARDS  HENDERSON 

PROFESSORS   OF  CHEMISTRY,  OHIO   STATE  UNIVERSITY 


GINN  AND  COMPANY,  BOSTON 

NEW  YORK  •  CHICAGO  •  LONDON 
ATLANTA  •  DALLAS  •  COLUMBUS  •  SAN  FRANCISCO 


GIFT 


COPYRIGHT,  1915,  BY 
WILLIAM  McPHERSON  AND  WILLIAM  E.  HENDERSON 


ALL  BIGHTS   RESERVED 
315.3 


DEFT. 


llf*  »Krr. 


GINN  AND  COMPANY  •  PRO 
PRIETORS  •  BOSTON  •  U.S.A. 


PREFACE 

In  preparing  this  introduction  to  the  science  of  chemis- 
try the  authors  have  endeavored  to  provide  a  text  easily 
within  the  grasp  of  the  average  high-school  student.  Their 
effort  has  been  to  make  the  subject  interesting,  to  use  sim- 
ple language,  to  develop  the  theoretical  portions  briefly  and 
as  a  natural  outcome  of  facts  already  presented,  and  to  em- 
phasize as  much  as  possible  the  applications  of  chemistry  in 
everyday  life.  The  authors  realize  that  the  great  majority 
of  the  students  of  chemistry  in  our  secondary  schools  will 
never  go  further  in  the  formal  study  of  the  science,  and 
the  book  is  primarily  for  them.  It  will  be  found,  however, 
that  the  requirements  of  the  College  Entrance  Examina- 
tion Board  have  been  fully  met. 

In  addition  to  the  applications  to  metallurgy  and  to 
manufacturing  which  are  always  of  interest,  an  unusual 
amount  of  matter  relating  to  agriculture,  to  household 
chemistry,  and  to  sanitation  has  been  introduced.  Since 
the  compounds  of  carbon  play  so  important  a  part  in  our 
daily  life,  their  discussion  has  been  made  more  ample  than  is 
usual  in  an  elementary  course,  and  the  chapters  devoted  to 
these  compounds  have  been  brought  into  their  proper  place 
in  the  text. 

While  thus  emphasizing  in  every  legitimate  way  the  appli- 
cations of  chemistry,  the  authors  have  never  lost  sight  of 
the  fact  that  they  are  introducing  young  minds  to  the  ele- 
ments of  a  wonderful  science  and  that  it  is  due  both  to  the 


vi  FIRST  COURSE  IX  CHEMISTRY 

student  and  to  the  science  to  exercise  care  in  diction,  candor 
in  spirit,  moderation  in  statement,  and  logic  in  presentation. 
It  would  not  be  a  difficult  task  to  write  a  more  interesting: 

o 

book  of  popular  chemical  information  —  a  sort  of  wonder- 
book  of  unconnected  facts ;  but  the  authors  feel  that  all 
students,  particularly  those  who  do  not  expect  to  continue 
their  work  in  college,  should  have  some  just  impression  of 
the  spirit  of  a  science  and  of  scientific  thought  as  well  as 
of  the  contributions  of  the  science  to  human  comfort.  Some 
responsibility  for  vivacity,  for  interesting  bits  of  collateral 
information,  and,  especially,  for  stimulating  interest  in  the 
practical  applications  of  chemistry  as  illustrated  in  local 
industries  must  always  rest  upon  the  teacher. 

Directions  for  laboratory  work  will  be  found  in  a  sepa- 
rate volume  entitled  "  Exercises  in  Chemistry."  While  these 
exercises  have  been  chosen  primarily  to  demonstrate  the 
principles  developed  in  the  textbook,  great  care  has  been 
exercised  to  reduce  the  requirements  for  apparatus  and 
the  cost  of  chemicals  to  a  minimum.  In  the  Appendix  will 
be  found  some  suggestions  in  regard  to  theme  writing.  The 
authors  are  convinced  that  many  advantages  will  be  gained 
by  following  up  these  suggestions. 

A  great  deal  of  labor  has  been  expended  in  securing  the 
illustrations  for  the  text,  and  it  is  believed  that  these  will 
add  much  to  the  interest  of  the  book  and  to  the  clearness  of 
the  presentation.  The  authors  gratefully  acknowledge  their 
indebtedness  to  a  number  of  their  colleagues,  especially  to 
James  R.  Withrow  and  John  F.  Lyman,  and  to  Robert  C. 
Hummell,  who  has  been  kind  enough  to  read  the  proof  sheets. 
They  are  also  under  obligations  to  many  other  individuals 
and  to  manufacturing  firms  and  wish  to  mention  especially  the 
following :  Charles  Hoover,  in  charge  of  the  water-filtration 


PREFACE  vii 

plant,  and  Clarence  Hoover,  in  charge  of  the  sewage-disposal 
plant,  Columbus,  Ohio;  Donald  (I.  Kohr  of  the  Lowe 
Brothers  Paint  Co. ;  Frank  O.  Clements  of  the  National 
Cash  Register  Co.  ;  Dr.  C.  H.  Viol  of  the  Standard  Chemi- 
cal Co. ;  Dr.  H.  P.  Armsby,  Pennsylvania  State  College ; 
R.  E.  Humphrey,  chief  chemist,  Standard  Oil  Co.;  The 
Pittsburgh  Plate  Glass  Co. ;  The  H.  L.  Dixoii  Co. ;  The 
American  Steel  and  Wire  Co. ;  The  Hydraulic  Press 
Manufacturing  Co. ;  The  German  Kali  Works ;  The  Chili 
Niter  Co. ;  The  Picher  Lead  Co.  ;  The  United  States  Beet 
Sugar  Industry  ;  The  Keever  Starch  Co. ;  The  Lackawanna 
Steel  Co. ;  The  American  Rolling  Mills  Co. 

THE  AUTHORS 

OHIO  STATI;   UXIVKKSITY 
COLUMBUS,  OHIO 


CONTENTS 

CHAPTER  PAGE 

I.   BURNING 1 

II.    ELEMENTS  AND  COMPOUNDS 8 

III.  OXYGEN 14 

IV.  HYDROGEN 24 

V.    THE  GAS  LAWS;    STANDARD  CONDITIONS  ....  33 

VI.    COMPOUNDS  OF  HYDROGEN  AND  OXYGEN  :    WATER 

AND  HYDROGEN  PEROXIDE 40 

VII.    MATTER  AND  ENERGY 55 

VIII.    COMBINING  WEIGHTS  ;    THE   ATOMIC  THEORY     .     .  65 

^     IX.    FORMULAS;    EQUATIONS;    CALCULATIONS    ....  71 
X.    NITROGEN    AND    THE    RARE    ELEMENTS:     ARGON, 

HELIUM,   NEON,   KRYPTON,   XENON 80 

XI.    THE  ATMOSPHERE 86 

XII.    SOLUTIONS  AND  TONIJSATION 96 

XIII.  ACIDS,  BASES,  AND  SALTS;    NEUTRALIZATION     .     .  107 

XIV.  VALENCE 116 

XV.    COMPOUNDS  OF  NITROGEN 121 

-^     XVI.    EQUILIBRIUM;    MASS  ACTION 135 

XVII.   SULFUR  AND  ITS  COMPOUNDS 140 

XVIII.    THE  PERIODIC  LAW 159 

XIX.    THE  CHLORINE  FAMILY     .     ,     , 166 

-?    XX.    MOLECULAR   WEIGHTS  ;    ATOMIC  WEIGHTS      .     .     .  185 

'       XXI.    CARBON  AND  SOME  OF  ITS  SIMPLER  COMPOUNDS    .  193 

XXII.    FUELS;    FLAMES;    ELECTRIC  FURNACES      ....  210 

XXIII.  CARBOHYDRATES  ;      ALCOHOLS  ;      COAL-TAR      COM- 

POUNDS   220 

XXIV.  ORGANIC  ACIDS  ;  FATS  ;  OILS  ;  PROTEINS    ....  235 
XXV.    FOODS 240 

XXVI.    THE  PHOSPHORUS  FAMILY  247 


x  FIRST  COURSE  IX  CHEMISTRY 

CHAPTER  PAGE 

XXVII.    SILICON  AND  BORON 260 

XXVIII.    THE  METALS 272 

XXIX.    THE  SODIUM  FAMILY 27(5 

XXX.    SOAP;    GLYCERIN;    EXPLOSIVES 291 

XXXI.    THE  CALCIUM  FAMILY 297 

XXXII.    FERTILIZERS 309 

XXXIII.  THE  MAGNESIUM  FAMILY     ........  313 

XXXIV.  ALUMINIUM 322 

XXXV.    ALUMINIUM  SILICATES  AND  THEIR  COMMERCIAL 

APPLICATIONS 333 

XXXVI.    THE  IRON  FAMILY 338 

XXXVII.    COPPER,  MERCURY,  AND  SILVER 357 

XXXVIII.    TIN  AND  LEAD 369 

XXXIX.    URANIUM  AND  RADIUM 379 

XL.    MANGANESE  AND  CHROMIUM 385 

XLI.    PLATINUM   AND  GOLD 391 

XLII.    SOME  APPLICATIONS  OF   RARER    KLEMEXTS    .     .  396 

APPENDIX 399 

THEME-WRITING 399 

LIST  OF  BOOKS  FOR  REFERENCE 400 

TENSION  OF  AQUEOUS  VAPOR 402 

WEIGHT  OF  1  LITER  AND  BOILING  POINT  OF  VARIOUS 

GASES 402 

DENSITIES  AND  MELTING  POINTS  OF  SOME  COMMON 

ELEMENTS 402 

SOLUBILITIES  OF  GASES  IN  WATER 403 

SOLUBILITIES  OF  SOLIDS 403 

RELATION  OF  COMMON  UNITS  AND  METRIC  UNITS  .  .  403 

INDEX 405 

PERIODIC   ARRANGEMENT  OF  THE   ELEMENTS 

Facing1  hack  cover 

LIST  OF   ELEMENTS                 Inside  back  cover 


FIRST  COURSE  IN  CHEMISTRY 


CHAPTER  I 
BURNING 

Introduction.  Very  few  of  us  can  watch  a  fire  raging 
without  mingled  feelings  of  wonder  and  awe.  We  are  im- 
pressed by  the  evidence  it  presents  of  uncontrolled  forces 
at  work ;  we  are  conscious  of  a  sense  of  mystery  as  to 
what  becomes  of  the  things  that  have  been  consumed ;. 
we  wonder  about  the  cause  of  the  spectacle.  We  are 
not  surprised  that  in  earlier  ages  fire  and  other  similar 
occurrences  —  lightning,  volcanic  eruptions,  the  aurora 
of  northern  skies  —  should  have  inspired  worship,  awak- 
ened .  fear,  aroused  curiosity,  and  at  length  led  to  patient 
and  laborious  study. 

That  the  explanation  of  burning  is  by  no  means  a 
simple  matter  is  shown  by  the  fact  that,  notwithstanding- 
all  the  study  which  had  been  devoted  to  it,  it  was  not 
until  about  the  time  that  our  national  war  for  independ- 
ence was  being  fought,  less  than  a  century  and  a  half 
ago,  that  the  fundamental  facts  of  combustion  were  recog- 
nized. It  will  be  instructive  to  follow  a  few  experiments 
attentively  and  see  what  these  facts  are. 

A  burning  candle.  Two  facts  about  a  burning  candle  are 
nt  once  apparent :  (1 )  the  material  of  which  the  candle  is 
composed  disappears;  and  (2)  heat  and  light  are  given  out 
in  the  process.  If  a  dry  wide-mouthed  bottle  is  inverted 

i 


FIRST  COURSE  IN  CHEMISTRY 


FIG.  1.  Collecting  the 

products  formed  from 

a  burning:  candle 


bv^V  the  Haine  (Fig.  1),  a  further  fact  can  be  learned : 

as  thV  candle  biirtis;\  moisture  is  deposited  upon  the  cold 
inside  walls  of  the  bottle.  Furthermore, 
if  the  bottle  is  removed  from  over  the 
flame  and  a  little  clear  limewater  at  once 
poured  into  it  and  gently  shaken,  the 
clear  liquid  becomes  milky,  whereas  it 
does  not  do  this  to  any  considerable 
extent  if  the  bottle  has  not  been  held 
over  the  flame.  It  is  therefore  evident 
that,  in  burning,  the  candle  is  not  really 
destroyed,  but  from  it  are  formed  water 
and  another  substance  which  renders 
clear  limewater  milky.  This  latter  sub- 
stance is  a  gas  and  is  known  as  carbon 
dioxide  or  carbonic  acid  gas. 
Increase  in  weight  during  burning.  If  the  experiment 

is  repeated  under  the  conditions  represented  in  Fig.  2,  an 

additional  fact  may  be  learned. 

The  candle  A  is  placed  on  one 

pan  of  the  balance.    Over  it 

is  suspended  a  wide  glass  tube 

B  (a  lamp  chimney  serves  very 

well)  loosely  filled  with  pieces 

of  quicklime  or  caustic  potash, 

both  of  which  substances  have 

been  found  to  absorb  moisture 

and  also  carbon  dioxide.    The 

whole  apparatus  is   carefully 

counterpoised  by  weights  (7; 

then  the  candle  is  lighted.   As 

the  candle  burns,  the  pan  upon  which  it  rests  gradually 

sinks,  indicating  that  the  gases  formed  during  burning  and 


FIG.  2.  Increase  in  weight  durin; 
burning 


BURKING 


3 


absorbed  in  the  tube  over  the  flame  are  heavier  than  the 
part  of  the  candle  which  has  burned. 

The  burning  of  iron.  Some  additional  information  may  be 
obtained  by  burning  iron.  Two  or  three  grams  of  fine  iron 
powder  is  placed  in  a  small  evaporating  dish  and  accu- 
rately weighed.  The  dish  containing  the  iron  is  then  strongly 
heated  by  a  laboratory  lamp  (known  as  a  Bunsen  burner), 
as  shown  in  Fig.  3.  As  the  heating  proceeds,  no  fumes 
or  gases  can  be  discovered,  but  the 
iron  glows  and  turns  into  a  dark-red 
powder  which  in  no  way  resembles 
the  original  iron.  If  the  crucible 
and  contents  are  now  cooled  and 
weighed  again,  they  will  be  found  to 
weigh  more  than  before  the  iron  was 
burned.  Other  metals,  such  as  lead, 
tin,  and  zinc,  act  in  a  similar  way,  the 
product  formed,  or  the  ash,  always 
being  heavier  than  the  unburned 
metal. 


This  increase  in  weight  is  one  of  FlG>3>  The  burning  of  iron 
the  most  striking  facts  of  burning. 

Unless  we  assume  that  weight  has  been  created  out  of 
nothing  (which  is  contrary  to  all  our  experience),  we  must 
suppose  that  substances,  in  burning,  withdraw  from  the  air 
something  possessing  weight.  This  suggestion  can  be  put 
to  a  test  at  once. 

The  burning  of  phosphorus.  A  piece  of  phosphorus  the 
size  of  a  pea  is  placed  upon  a  crucible  lid  supported  on  a 
flat  cork  floating  in  a  trough  of  water  (Fig.  4).  The  phos- 
phorus is  set  on  fire  by  touching  it  with  a  lighted  match, 
and  a  good-sized  bottle  or  bell  jar  is  inverted  over  it,  as  is 
shown  in  the  figure.  At  first  a  few  bubbles  of  air  are  forced 


FIRST  COUESE  IN  CHEMISTRY 


FIG.  4.    The  burning  of 
phosphorus 


out,  owing  to  the  expansion  of  the  air  occasioned  by  the  heat. 
Then  the  water  rises  in  the  bell  jar,  showing  that  air  is  being 
used  up.  Soon  the  phosphorus  ceases  to  burn,  although  there 
is  plenty  of  air  left.  Very  careful  experiments  show  that 
burning  will  continue  until  about 
one  fifth  of  the  air  has  been  used 
up,  and  will  then  cease.  The  smoke 
of  the  burning  phosphorus  grad- 
ually disappears,  owing  to  the  fact 
that  it  has  dissolved  in  the  water. 

This  experiment  suggests  that  some 
material,  constituting  about  one  fifth 
of  the  air  by  volume,  is  withdrawn 
from  the  air  by  a  burning  substance  such  as  phosphorus. 
Experiment  of  Priestley.  The  clue  to  the  nature  of  the 
material  withdrawn  from  the  air  was  first  found  by  Joseph 
Priestley,  an  Englishman,  in 
1774.  He  obtained  a  red  solid 
(which  was  then  called  red  pre- 
cipitate of  mercury,  but  which 
is  merely  the  ash  of  mercury 
burned  in  the  air)  and  heated  it 
strongly  in  a  closed  tube.  He 
found  that  a  remarkable  gas  (or 
air,  as  he  called  it)  was  given  off, 
in  which  a  candle  burned  much 
more  brilliantly  than  in  ordinary 
air,  while  mercury  was  left  in 
the  tube.  The  experiment  may  be  repeated  as  is  shown  in 
Fig.  5,  introducing  a  splinter  of  wood  with  a  glowing  spark 
on  the  end  into  the  tube  in  which  the  red  precipitate  is 
being  heated.  The  spark  at  once  bursts  into  a  flame.  It 
was  later  shown  that  phosphorus  will  use  up  all  of  this  air 


FIG. 


Heating  the  ash  of 
mercury 


BUENING 


in  burning.  Priestley  did  not  realize  the  importance  of  his 
experiment,  but  its  significance  was  quickly  understood  by 
the  French  chemist  Lavoisier  (see  frontispiece). 

Experiment  of  Lavoisier.  Lavoisier  believed  that  air  is 
in  part  composed  of  this  gas,  which  he  named  oxygen,  and 
that,  in  burning,  substances  combine  with  oxygen.  To  test 
this  idea  he  devised  the  following  experiment : 

Details  of  Lavoisier's  experiment.  A  weighed  quantity  of  pure 
mercury  was  placed  in  a  retort  A  (Fig.  6),  which  communicated 
with  a  bell  jar  B con- 
taining air  confined 
over  mercury  in  C. 
The  retort  and  its 
contents  were  heated 
by  the  furnace  D 
for  twelve  days  just 
below  the  boiling 
point  of  mercury, 
and  it  was  found 
that  a  part  of  the 
mercury  had  been 


B 


FIG.  6.    Apparatus  used  by  Lavoisier  to  prove 
that  mercury  absorbs  oxygen  during  burning 


burned,  forming  a 
red  ash,  and  that 
from  7  to  8  cu.  in.  of  air  had  been  absorbed  during  the  process. 
The  red  ash  was  collected  and  placed  in  a  small  retort  arranged 
in  such  a  way  that  any  gas  given  off  from  it  could  be  collected 
and  measured.  It  was  then  strongly  heated,  by  which  process 
the  mercury  was  recovered  and  a  gas  collected.  In  quantity 
this  proved  to  be  equal  to  from  7  to  8  cu.  in.,  or  the  same  vol- 
ume that  had  been  previously  absorbed.  It  was  not  ordinary 
air,  however,  but  oxygen.  In  this  way  a  definite  quantity  of 
mercury  was  first  burned  in  air,  and  the  ash  formed  was  then 
decomposed  into  mercury  and  oxygen. 

Conclusions.    These    experiments,    together   with  many 
later  and  much   more   accurate   ones,  have  demonstrated 


6  FIRST  COURSE  IN  CHEMISTRY 

that  burning  is  really  a  union  between  the  material  burned 
and  the  oxygen  which  is  present  in  the  air,  and  that  the 
total  weight  of  the  material  burned  plus  that  of  the  oxygen 
used  up  equals  the  weight  of  the  products  formed.  It  will 
be  seen  that  the  discovery  of  oxygen  and  the  demonstration 
of  the  real  nature  of  burning  were  greatly  aided  by  the 
happy  chance  that  when  the  ash  of  mercury  is  strongly 
heated,  the  act  of  burning  is  turned  around,  or  reversed, 
oxygen  and  mercury  being  recovered.  This  fact  may  be 
represented  in  an  equation  as  follows: 

mercury  ash  <    >  mercury  +  oxygen 

The  double  arrows  indicate  that  the  change  may  take  place 
in  either  direction,  according  to  the  conditions. 

Additional  questions.  This  still  leaves  many  questions 
unanswered.  What  causes  the  combustion  to  take  place  ? 
Why  will  some  things  burn  and  others  not?  Where  do 
the  heat  and  light  come  from,  and  what  finally  becomes 
of  them  ?  In  part  these  questions  will  be  answered  as 
we  go  along. 

The  phlogiston  theory  of  burning.  The  reason  why  Priestley 
failed  to  understand  the  part  which  oxygen  plays  in  burning 
was  that  he  shared  the  prevailing  views  of  his  time  and  could 
not  easily  give  up  the  ideas  he  had  always  held.  For  a  century 
and  a  half  burning  had  been  regarded  as  occasioned  by  the 
escape  of  an  invisible  material  known  as  phlogiston.  Doubt- 
less this  idea  was  suggested  by  the  puffing  and  sputtering  of 
materials  while  they  burn.  It  was  thought  that  when  mercury 
burns,  a  constituent  of  it,  namely,  phlogiston,  escapes,  so  that 
the  ash  formed  is  equal  to  the  mercury  minus  phlogiston.  To 
recover  the  mercury  from  the  ash,  phlogiston  must  be  restored 
to  it,  and  it  was  hard  to  see  how  merely  heating  it  could  do 
this.  Priestley  thought  that  the  oxygen  given  off  must  be  air 
from  which  phlogiston  had  been  extracted  and  returned  to  the 


BURNING  7 

mercury,  and  he  called  it  dephlogisticated  air.  Evidently  this 
theory  could  not  explain  why  metals  get  heavier  as  they  burn, 
and  when  oxygen  was  discovered,  the  theory  of  phlogiston  was 
soon  abandoned. 

EXERCISES 

1.  How  do  the  ashes  of  coal  compare  in  weight  with  the  original 
coal  ?    Can  you  suggest  a  reason  why  some  coals,  when  burned,  leave 
more  ashes  than  others  ? 

2.  How  do  you  account  for  the  fact  that  when  a  lamp  is  lighted, 
a  film  of  moisture  is  deposited  on  the  chimney  ?    Is  the  film  deposited 
on  the  inside  or  on  the  outside  of  the  chimney  ?  Why  does  the  film 
soon  disappear? 

3.  Could  you  improve  the  details  of  the  apparatus  illustrated  in 
Fig.  4  so  that  no  air  would  be  forced  out  ? 

4.  How  accurately  does  Lavoisier  seem  to  have  measured  his 
gases  ?    What  is  the  percentage  error  ?   Do  you  think  his  conclusions 
from  such  experiments  were  justified  ? 

5.  How  did  the  fact  that  the  burning  of  mercury  is  reversible 
aid  in  the  discovery  of  oxygen  ? 

6.  In  Fig.  6  how  would  it  be  known  that  air  had  been  absorbed  ? 

7.  It  was  known  to  the  phlogistonists  that  things  would  not  burn 
in  the  absence  of  air.    Was  this  in  accord  with  their  theory  ? 

8.  What  is  the  derivation  of  the  word  phlogiston  (see  dictionary)? 

TOPICS  FOR  THEMES  * 

Lavoisier  (Thorpe,  Essays  in  Historical  Chemistry). 
Phlogiston  (see  encyclopedia). 

*  Refer  to  Appendix  for  suggestions  in  connection  with  topics  for  themes. 


CHAPTER  II 
ELEMENTS  AND  COMPOUNDS 

Elements.  Having  found  that  by  merely  heating  the  ash 
of  mercury  (or  mercuric  oxide,  as  it  is  now  called)  two  en- 
tirely different  materials,  mercury  and  oxygen,  are  obtained 
from  it,  the  inquiry  naturally  arises :  Can  still  other  sub- 
stances be  obtained  by  heating  either  of  these  two  ?  Will 
any  agency  other  than  heat  bring  about  their  decomposi- 
tion ?  Many  efforts  to  decompose  these  substances  have 
been  made,  but  none  have  succeeded.  On  this  account  they 
are  called  elementary  substances  or  elements. 

It  is  not  always  easy  to  prove  that  a  given  substance  is 
really  an  element.  Some  way  as  yet  untried  may  be  suc- 
cessful in  decomposing  it  into  other  simpler  forms  of  matter. 
Water,  lime,  and  many  other  familiar  substances  were  at 
one  time  thought  to  be  elements,  but  are  now  known  to 
contain  two  or  more  elements. 

Compounds.  Substances  such  as  mercuric  oxide,  which 
are  formed  by  the  union  of  two  or  more  elements,  are  called 
compounds.  Experiments  have  shown  that  any  given  com- 
pound is  always  made  up  of  the  same  elements  and  that  it 
has  a  perfectly  definite  percentage  of  each  one  ;  for  example, 
mercuric  oxide  is  always  composed  of  mercury  and  oxygen 
combined  in  the  proportion  92.6  per  cent  mercury  and  7.4 
per  cent  oxygen.  We  shall  learn  of  other  characteristics 
of  compounds  as  we  proceed. 

As  a  rule  the  appearance  of  a  compound  offers  no  sug- 
gestion as  to  what  elements  are  present  in  it.  Thus,  the 

8 


ELEMENTS  AND  COMPOUNDS  9 

red  solid,  mercuric  oxide,  is  formed  by  the  union  of  the  sil- 
very liquid,  mercury,  with  the  invisible  gas,  oxygen.  The 
familiar  colorless  liquid,  water,  is  formed  by  the  union  of 
two  invisible  gases,  oxygen  and  hydrogen. 

Chemical  changes  ;  chemical  action.  Such  changes  as 
those  taking  place  in  the  conversion  of  mercuric  oxide  into 
mercury  and  oxygen  or  in  the  burning  of  any  substance 
are  called  chemical  changes,  and  in  describing  them  we  say 
that  chemical  action  has  taken  place.  In  all  such  changes 
the  substances  resulting  from  the  chemical  action  differ  in 
composition  from  the  substances  originally  present,  and 
usually  differ  from  them  in  appearance  as  well.  We  shall 
see  later  on  that  there  are  other  important  changes  which 
always  accompany  chemical  action. 

Chemical  affinity.  The  force  that  causes  elements  to 
unite  and  holds  them  in  combination  in  compounds  is 
called  chemical  affinity.  We  know  very  little  about  the 
nature  of  this  force,  just  as  we  know  very  little  about  the 
force  of  gravitation.  It  is  evident,  however,  that  there  is 
such  a  force,  and  it  is  convenient  to  have  a  name  by  which 
we  can  refer  to  it. 

Number  of  elements.  The  number  of  substances  now 
considered  to  be  elements  is  not  large  —  about  eighty  in 
all.  Many  of  these  are  rare,  and  very  few  of  them  form  any 
large  fraction  of  the  materials  in  the  earth's  crust.  Clarke 
gives  the  following  estimate  of  the  composition  of  the  solid 
portion  of  the  earth's  crust : 

COMPOSITION  OF  THE  EARTH'S  CRUST 

Oxygen 47.07%  Magnesium     ....  2.40% 

Silicon 28.06%  Sodium 2.43% 

Aluminium  .     .     .     .  7.90%  Potassium 2.45% 

Iron 4.43%  Hydrogen 0.22% 

Calcium 3.44%  Other  elements  ,  l.( 


10  FIRST  COURSE  IN  CHEMISTRY 

A  complete  list  of  the  elements  is  given  on  the  back 
cover  page.  In  this  list  the  more  common  of  the  elements 
are  printed  in  heavier  type.  It  is  not  necessary  to  study 
more  than  one  third  of  the  total  number  of  elements  to 
gain  a  very  good  knowledge  of  chemistry. 

Elements  in  the  human  body.  Comparatively  few  of  the 
elements  appear  to  be  essential  to  life.  The  following  table, 
compiled  by  Sherman,  gives  the  average  composition  of  the 
human  body.  So  far  as  we  can  judge,  these  are  the  only 
ones  upon  which  living  organisms  are  dependent,  though 
traces  of  others  may  be  necessary. 

AVERAGE  COMPOSITION  OF  THE  HUMAN  BODY 

Oxygen  .  65.00%  Phosphorus  1.00%  Magnesium  .  0.05% 

Carbon    .  18.00%  Potassium  0.35%  Iron     .     .     .  0.004% 

Hydrogen  10.00%  Sulfur  .     .  0.25%  Iodine      .     .  traces 

Nitrogen  3.00%  Sodium      .  0.15%  Fluorine  .     .  traces 

Calcium .  2.00%  Chlorine    .  0.15%  Silicon      .     .  traces 

Occurrence  of  the  elements.  Most  of  the  elements  occur 
in  nature  not  as  un combined  substances,  but  in  the  form  of 
chemical  compounds.  When  an  element  does  occur  uncom- 
bined,  as  is  the  case  with  gold  and  sulfur,  we  say  that  it 
occurs  in  the  free  state  or  native ;  when  it  is  combined  with 
other  substances  in  the  form  of  compounds,  we  say  that  it 
occurs  in  the  combined  state,  or  in  combination. 

Names  of  elements.  The  names  given  to  the  elements 
have  been  selected  in  a  great  many  different  ways.  Some 
names,  such  as  iron  and  gold,  are  very  old,  and  their  original 
meaning  is  obscure.  Many  names  indicate  some  striking 
property  of  the  element.  The  name  bromine,  for  example, 
means  "  stench,"  referring  to  the  extremely  unpleasant  odor 
of  the  substance.  Other  elements  are  named  from  countries 
or  localities,  as  germanium  and  scandium. 


ELEMENTS  AND  COMPOUNDS 


11 


Symbols.  In  indicating  the  elements,  chemists  have 
adopted  a  system  of  abbreviations.  These  are  known  as 
symbols,  each  element  having  a  distinctive  symbol.  Some- 
times the  initial  letter  of  the  name  is  adopted  to  indicate  the 
element.  Thus,  I  stands  for  iodine,  C  for  carbon.  Usually 
it  is  necessary  to  add  some  other  characteristic  letter  to  the 


FIG.  7.   An  alchemist's  laboratory,  in  the  Deutsches  Museum,  Munich 


symbol,  since  several  names  may  begin  with  the  same  letter. 
Thus,  C  stands  for  carbon,  Cl  for  chlorine,  Cd  for  cadmium. 
Sometimes  the  symbol  is  an  abbreviation  of  the  name  in 
some  other  language.  In  this  way  Fe  (Latin,  ferruni)  in- 
dicates iron.  The  symbols  will  be  found  in  the  list  of  ele- 
ments given  on  the  back  cover  page.  They  will  become 
familiar  through  constant  use. 


12  FIRST  COURSE  IN  CHEMISTRY 

The  number  of  compounds.  The  number  of  compounds 
which  have  been  described  and  which  can  be  made  when 
desired  is  very  large,  and  each  year  many  more  are  added 
to  the  list.  About  200,000  are  known  that  contain  the  ele- 
ment carbon  as  one  constituent,  and  the  total  number  listed 
in  the  large  handbooks  of  chemistry  is  much  larger.  Fortu- 
nately it  is  not  necessary  to  become  familiar  with  any  large 
number  of  these  in  order  to  gain  an  understanding  of  the 
principles  of  chemistry. 

Alchemy.  In  olden  times  it  was  thought  that  some  way  could 
be  found  to  change  one  element  into  another,  and  a  great  many 
experiments  were  made  to  accomplish  this  transformation. 
Most  of  these  efforts  were  directed  toward  changing  the  com- 
moner metals  into  gold,  and  many  fanciful  ways  for  doing  this 
were  described.  The  chemists  of  that  time  were  called  alche- 
mists, and  the  art  which  they  practiced  was  called  alchemy. 
Failing  to  accomplish  this  transformation,  the  alchemists  grad- 
ually became  convinced  that  the  only  way  in  which  common 
metals  could  be  changed  into  gold  was  by  the  wonderful  power 
of  a  magic  substance  which  they  called  the  philosophers  stone, 
which  would  accomplish  this  transformation  by  its  mere  touch 
and  would  in  addition  give  perpetual  youth  to  its  fortunate 
possessor.  No  one  has  ever  found  such  a  stone,  but  one  of  the 
most  brilliant  discoveries  of  modern  times  has  demonstrated  that 
at  least  some  of  the  elements  are  of  their  own  accord  very  slowly 
changing  into  others. 

EXERCISES 

1.  What  other  means  of  decomposing  a  compound  can  you  sug- 
gest, aside  from  heating? 

2.  How  would  you  define  a  compound ?   an  element? 

3.  What  is  meant  by  the  earth's  crust? 

4.  Does  the  fact  that  a  substance  undergoes  no  change  on  heat- 
ing show  it  to  be  an  element  ? 

5.  Read  over  the  list  of  elements.    What  ones  do  you  know  to 
occur  native? 


ELEMENTS  AND  COMPOUNDS  13 

*6.    Aluminium  is  much  more  abundant  than  iron  (see  table). 
How  do  you  account  for  the  much  greater  cheapness  of  iron? 

7.  Consult  the  dictionary  for  the  derivation  and  significance  of 
the  following  names :  phosphorus,  hydrogen,  germanium,  columbium, 
chlorine,  argon,  copper,  selenium,  thorium. 

8.  How  would  you  define  (a)  chemical  changes  and  (ft)  chemical 
action  ? 

9.  What  weight  of   oxygen  is  present  in    500  g.  of   mercuric 
oxide?    Ans.  37  g. 

TOPICS  FOR  THEMES 

Methods  of  naming  the  elements,  with  examples  to  illustrate  (see 

dictionary). 

The  philosopher's  stone  and  the  elixir  of  life  (see  encyclopedia). 
The  alchemists   (Muir,  The  Story  of  Alchemy;    Bird,  Modern 

Science  Reader). 


CHAPTER  III 
OXYGEN 

Introduction.  Having  become  acquainted  with  a  few  of 
the  characteristics  of  the  class  of  substances  called  elements, 
we  shall  now  turn  to  a  more  detailed  study  of  two  members 
of  this  class,  namely,  oxygen  and  hydrogen.  It  is  natural 
that  we  should  begin  with  oxygen,  since  it  is  the  most 
abundant  of  all  elements,  occurs  in  nature  in  great  quanti- 
ties in  the  elementary  state,  and  plays  such  an  important 
part  in  the  familiar  processes  of  burning  and  breathing. 

Discovery.  Priestley's  experiment  (1774)  of  heating 
mercuric  oxide  (p.  4)  is  looked  upon  as  constituting  the 
discovery  of  oxygen,  though  it  is  now  known  that  the 
Swedish  chemist,  Scheele,  had  prepared  it  some  years  earlier 
by  heating  niter.  An  account  of  this  latter  experiment  was 
not  published  until  1777,  while  Priestley  at  once  made 
known  the  results  of  his  experiment. 

The  name  oxygen.  The  name  oxygen,  suggested  by  Lavoisier, 
means  "acid  producer/'  for  he  thought  that  the  class  of  sub- 
stances known  as  acids  owe  their  characteristic  properties  to 
the  presence  in  them  of  this  element.  We  now  know  that 
there  are  acids  which  contain  no  oxygen. 

Occurrence.  Oxygen  is  by  far  the  most  abundant  of  the 
elements.  In  the  free  state  it  forms  a  considerable  part  of 
the  atmosphere,  100  volumes  of  dry  air  containing  about 
21  volumes  of  oxygen.  Combined  with  other  elements,  it 

14 


OXYGEN 


15 


forms  eight  ninths  of  water,  nearly  one  half  of  the  rocks 
constituting  the  earth's  crust,  and  over  one  half  of  ani- 
mal and  vegetable  organisms;  for  example,  65  per  cent 
by  weight  of  the  human  body  is  oxygen. 

Preparation.  Among  the  great  number  of  compounds 
containing  oxygen  there  are  a  few  which  can  be  decom- 
posed easily  in  such  a  way  as  to  set  the  oxygen  free. 

1.  Preparation  from  mercuric  oxide.    Mercuric  oxide  has 
been  found  to  consist  of  7.4 

per  cent  oxygen  and  92.6  per 
cent  mercury.  If  a  small 
quantity  of  the  red  powder 
is  placed  in  a  narrow  test 
tube  and  heated  in  a  Bunsen 
flame  (Fig.  5),  it  is  rapidly 
decomposed  into  its  constit- 
uent elements.  The  mer- 
cury is  seen  to  deposit  on 
the  sides  of  the  tube,  while 
the  presence  of  oxygen  is 
shown  by  the  fact  that  a 
glowing  spark  on  the  end 
of  a  splinter  of  wood  in- 
serted into  the  tube  bursts 
into  a  bright  flame.  The 
method  is  of  interest  because  of  its  simplicity  and  because 
it  first  led  to  the  discovery  of  oxygen.  It  is  too  expensive 
to  serve  as  a  laboratory  method. 

2.  Preparation  from  potassium  chlorate  (usual  laboratory 
method).   Potassium  chlorate  is  a  white  solid  which  has  been 
found  to  consist  of  31.9  per  cent  potassium,  28.9  per  cent 
chlorine,  and  39.2  per  cent  oxygen.    When  this  material 
is  heated  above  its  melting  point,  the  oxygen  is  given  off, 


FIG.  8.   Joseph  Priestley 
(1733-1804) 

The  discoverer  of  oxygeii 


16 


FIKST  COURSE  IN  CHEMISTKY 


leaving  a  compound  of  chlorine  and  potassium  called  potas- 
sium chloride.    The  changes  may  be  represented  as  follows : 


potassium  chlorate 

Tpotassiumn 
I  chlorine 
Loxygen 


potassium  chloride  4-  oxygen 

["potassium"! 
I  chlorine 


The  evolution  of  the  oxygen  begins  at  about  400°.  It 
has  been  found,  however,  that  if  a  small  quantity  of  cer- 
tain finely  powdered  solids,  such  as  manganese  dioxide,  is 
mixed  with  the  potassium  chlorate, 
not  only  is  the  oxygen  given  off  more 
uniformly,  but  the  rate  at  which  it 
is  evolved  at  any  given  temperature  is 
greatly  increased.  Just  how  the  man- 
ganese dioxide 
brings  about  this 
result  is  not  def- 
initely known. 
The  amount  of 
oxygen  obtained 
from  a  given 
weight  of  potas- 
sium chlorate  is  just  the  same  whether  the  manganese 
dioxide  is  present  or  not,  and  just  as  much  of  the  dioxide 
remains  at  the  end  as  was  added.  As  we  shall  see,  this 
kind  of  action  is  not  at  all  uncommon. 

Directions  for  preparing  oxygen.  A  convenient  way  of  pre- 
paring oxygen  from  potassium  chlorate  is  illustrated  in  the  ac- 
companying diagram  (Fig.  9).  A  mixture  consisting  of  four 
parts  of  potassium  chlorate  and  one  part  of  manganese  dioxide 
is  placed  in  the  flask  A  and  gently  heated.  The  oxygen  is 
evolved  and  escapes  through  the  tube  B.  It  is  collected  by 
bringing  over  the  end  of  the  delivery  tube  the  mouth  of  a 


FIG.  9.  The  preparation  of  oxygen  from  potassium 
chlorate 


OXYGEN 


17 


bottle  or  cylinder  C  completely  filled  with  water  and  inverted 
in  a  vessel  of  water  as  shown  in  the  figure.  The  gas  rises  in 
the  bottle  and  displaces  the  water. 

The  collection  of  gases.  The  method  just  described  for  col- 
lecting oxygen  illustrates  the  general  way  in  which  gases  are 
transferred  from  one  vessel  to  another  when  they  are  insoluble 
in  water  or  nearly  so.  The  vessel  D  (Fig.  9),  containing  the 
water  in  which  the  bottles  are  inverted,  is  called  a  pneumatic 
trim y1t.  Gases  which  are  soluble  in 
water  may  be  collected  in  a  similar 
way  over  mercury. 

3.  Commercial  preparation  of  oxygen. 

When  oxygen  is  prepared  for  com- 
mercial purposes,  it  is  usually  ob- 
tained either  from  water  or  from 
air.  As  we  shall  soon  see,  water 
contains  88.81  per  cent  of  oxygen 
and  11.19  per  cent  of  hydrogen.  It 
is  not  practicable  to  decompose  it 
into  its  elements  by  heat,  but  the 
decomposition  is  easily  effected  by 
the  use  of  electrical  energy.  The  FlG- 10«  The  decomposi- 

^  tion  of  water  into  oxygen 

method  may  be  illustrated  by  the  and  hydrogen  by  the  elec- 
following  laboratory  experiment :  trie  current 

Two  tubes,  A  and  B  (Fig.  10),  are  filled  with  water  and  in- 
verted in  a  vessel  of  water  to  which  a  little  sulfuric  acid  has 
been  added.  A  piece  of  platinum  foil,  C  and  D,  attached  to  a 
wire  is  then  brought  under  the  end  of  each  tube.  When  these 
wires  are  connected  with  a  suitable  source  of  current,  supply- 
ing from  6  to  10  volts  (about  6  dichromate  cells  in  series), 
bubbles  of  gas  will  be  seen  to  collect  in  each  tube.  These  gases 
are  oxygen  and  hydrogen.  The  volume  of  the  hydrogen  liberated 
is  twice  that  of  the  oxygen.  The  reasons  for  adding  sulfuric 
acid  will  be  explained  later  on. 


18  FIEST  COUESE  IN  CHEMISTRY 

This  process,  carried  out  on.  a  large  scale,  serves  as  a 
commercial  method  for  preparing  oxygen  and  hydrogen  as 
well.  Oxygen  is  also  obtained  commercially  from  air,  as  will 
be  described  later  (p.  93).  The  gas  so  prepared  is  pumped 
into  steel  tubes  (Fig.  11)  and  is  sold  in  this  form. 

Laboratory  methods  and  commercial  methods.  As  we  go  along 
we  shall  see  that  the  methods  used  in  making  various  substances 
in  the  laboratory  are  usually  different  from  those  employed  com- 
mercially. In  the  laboratory,  where  relatively  small  quantities 


FIG.  11.    Oxygen  ready  for  the  market 

are  desired,  the  easiest  or  most  instructive  way  is  preferred. 
In  commerce,  economy  is  the  deciding  factor.  Moreover,  it 
often  happens  that  a  method  which  will  not  work  well  on  a 
small  scale  works  admirably  with  commercial  quantities,  or 
that  the  value  of  a  second  product  (by-product^  obtained  at 
the  same  time  makes  a  method  a  success. 

Properties  of  a  substance.  By  the  properties  of  a  sub- 
stance we  mean  all  of  those  characteristics,  or  marks,  by 
which  we  recognize  it.  Some  of  these  are  concerned  with 
the  way  the  substance  affects  us,  such  as  taste,  odor,  color, 
and  luster.  Others  are  measurable  quantities,  such  as  den- 
sity, hardness,  solubility,  boiling  point,  and  freezing  point. 


OXYGEN 


19 


Properties  of  oxygen.  If  we  make  a  study  of  the  proper- 
ties of  oxygen,  we  find  that  it  is  a  colorless,  tasteless,  odor- 
less gas  slightly  heavier  than  air.  At  a  temperature  of  0° 
and  under  a  pressure  of  1  atmosphere,  1  liter  of  it  weighs 
1.4290  g.,  while  under  the  same  conditions  1  liter  of  air 
weighs  1.2928  g.  It  is  but  slightly  soluble  in  water.  Like 
other  gases,  oxygen  may  be  liquefied  by  applying  sufficient 
pressure  to  the  very  cold  gas.  This  liquid  is  pale  blue  in 
color.  It  boils  at  —  182.9°  and  freezes 
to  a  snowlike  solid  at  —  235°. 

Chemical  conduct.  In  addition  to  the 
properties  of  a  substance  we  always 
wish  to  know  how  it  conducts  itself 
toward  a  number  of  other  familiar  sub- 
stances, and  we  speak  of  this  as  the 
chemical  conduct  of  the  substance. 

At  ordinary  temperatures  oxygen  is 
not  very  active.  Most  substances  are 
either  not  affected  at  all  by  it  or  the 
action  is  so  slow  as  to  escape  notice. 
At  higher  temperatures,  however,  it  is 
very  active  and  unites  directly  with 
most  of  the  elements.  This  may  be  shown  by  heating 
various  substances  until  they  are  just  ignited  in  air,  and 
then  bringing  them  into  vessels  containing  oxygen,  when 
they  burn  with  greatly  increased  brilliancy.  Thus,  a  glow- 
ing splint  introduced  into  a  jar  of  oxygen  bursts  into  flame. 
Sulfur  burns  in  air  with  a  very  weak  flame  and  feeble  light ; 
in  oxygen  the  flame  is  increased  in  size  and  brightness 
(Fig.  12).  Substances  which  burn  readily  in  the  air,  such 
as  phosphorus,  burn  in  oxygen  with  dazzling  brilliancy. 
Even  substances  which  burn  in  the  air  with  great  difficulty, 
such  as  iron,  burn  readily  in  oxygen. 


FIG.  12.    Burning  sul- 
fur in  oxygen 


20  FIRST  COURSE  IN  CHEMISTRY 

Oxidation.  We  have  seen  that  the  burning  of  an  element 
in  the  air  is  really  its  union  with  oxygen.  Many  compounds 
as  well  as  elements  burn  readily  both  in  air  and  in  oxygen  ; 
among  these  are  coal,  wood,  oil,  and  gas.  In  the  majority 
of  such  cases  the  compound  is  completely  decomposed  and 
each  of  its  constituent  elements  combines  with  oxygen. 
Thus,  most  oils  are  made  up  of  carbon  and  hydrogen,  and 
when  the  oil  burns  it  is  converted  into  a  compound  of  car- 
bon and  oxygen  (carbon  dioxide)  and  another  of  hydrogen 
and  oxygen  (water).  Less  frequently  the  compound  under- 
goes no  decomposition  but  merely  as  a  whole  combines  with 
oxygen.  All  of  these  actions  are  called  oxidation.  W-e  shall 
see  later  that  still  other  kinds  of  chemical  action  are  called 
oxidation. 

Oxides.  Any  compound  consisting  of  oxygen  and  some 
one  other  element  is  called  an  oxide.  Thus,  burned  iron  is 
iron  oxide ;  burned  mercury  is  mercury  oxide ;  burned  oil 
yields  carbon  dioxide  and  hydrogen  oxide  (water).  When 
more  than  one  element,  aside  from  oxygen,  is  present  in  a 
compound,  it  is  not  called  an  oxide,  but  some  other  name 
is  given  it. 

Many  of  the  oxides  are  familiar  substances.  Thus,  water 
is  oxide  of  hydrogen  and  lime  is  oxide  of  calcium.  All  but 
about  half  a  dozen  elements  form  oxides,  and  many  of  them 
form  more  than  one,  so  that  a  large  number  of  these  com- 
pounds is  known.  Some  of  them  are  solid  bodies,  as  in  the 
case  of  the  oxides  of  mercury,  iron,  and  phosphorus  ;  others 
are  liquids,  of  which  class  water  is  the  most  familiar  exam- 
ple ;  quite  a  number  are  gases,  as  is  true  of  the  oxide  of 
carbon  and  of  sulfur. 

Combustion.  Sometimes  oxidation  takes  place  so  slowly 
that  no  light  is  seen,  and  unless  careful  measurements  are 
made,  no  heat  is  noticed.  The  decay  of  vegetable  matter 


OXYGEN  21 

such  as  wood  and  leaves,  is  an  example  of  this  slow  oxidation. 
In  other  cases,  as  with  burning  phosphorus  or  iron,  light  is 
given  off  either  as  a  flame  or  as  a  glow  called  incandescence. 
Oxidation  accompanied  by  light  is  called  combustion. 

Heat  of  oxidation  and  combustion.  Evidently  a  given 
substance  may  either  undergo  a  slow  oxidation  or  it  may 
undergo  combustion.  Thus,  a  piece  of  phosphorus,  exposed 
to  the  air  in  a  cold  room,  slowly  wastes  away  until  it  has 
all  disappeared  into  smoke  consisting  of  an  oxide  of  phos- 
phorus ;  but  if  it  is  touched  with  a  lighted  match,  it  takes 
fire  and  burns  very  rapidly,  giving  out  much  heat  in  its 
combustion.  The  product  is  the  same  in  both  cases,  namely, 
oxide  of  phosphorus.  Apparently  the  difference  lies  in  the 
amount  of  heat  given  off,  but  very  accurate  experiments 
demonstrate  that  this,  too,  is  exactly  the  same.  In  the  one 
case  the  action  is  so  slow  that  the  heat  is  conducted  away 
as  fast  as  it  is  liberated,  and  so  it  escapes  notice  ;  in  the  other 
it  is  given  off  so  rapidly  as  to  be  very  striking.  A  similar 
relation  has  been  found  to  hold  true  in  all  cases  of  combus- 
tion. The  heat  given  off  is  exactly  the  same  whether  the 
action  is  fast  or  slow,  provided  the  same  compound  is  formed. 

Spontaneous  combustion.  It  has  been  found  that  the  rate  at 
which  oxidation  goes  on  is  greatly  increased  by  raising  the  tem- 
perature of  the  material  undergoing  oxidation.  Consequently, 
if  the  conditions  surrounding  oxidation  are  such  that  the  heat 
given  off  cannot  escape,  the  temperature  will  steadily  rise,  and 
because  of  this  the  rate  of  oxidation  will  increase.  The  increased 
heat  thus  set  free  will  still  further  raise  the  temperature,  until 
the  oxidation  passes  into  active  combustion,  the  point  at  which 
this  occurs  being  called  the  kindling  temperature.  Materials  tak- 
ing fire  in  this  way  are  said  to  undergo  spontaneous  combustion. 
It  will  be  seen  that  the  essential  conditions  are  (1)  an  existing 
slow  oxidation  and  (2)  good  heat  insulation.  Linseed  oil,  used 
in  paints,  undergoes  rather  rapid  oxidation  in  air,  and  oily 


22  FIRST  COURSE  IN  CHEMISTRY 

rags  left  by  painters  not  infrequently  occasion  disastrous  fires. 
Fine,  dry  coal  in  the  center  of  a  heap  or  in  the  closed  hold  of 
a  vessel  sometimes  takes  fire.  Almost  any  finely  divided  com- 
bustible material,  such  as  sawdust  or  flour,  is  dangerous  when 
stored  in  a  warm,  dry  place.  Sometimes  the  heat  of  fermenta- 
tion, which  is  a  kind  of  oxidation,  will  start  a  fire  in  a  haystack 
or  barn  if  the  hay  is  not  well  dried  before  storing. 

Importance  of  oxygen.  Oxygen  is  one  of  the  most  im- 
portant of  the  elements.  It  is  essential  to  all  forms  of  life 
except  certain  low  forms  of  plant  life.  In  the  presence  of 


FIG.  13.    Sewage-disposal  plant,  Columbus,  Ohio,  in  which  the  sewage  is 
sprayed  into  the  air  to  secure  its  oxidation 

certain  minute  microorganisms,  which  in  some  way  assist 
in  the  process,  the  oxygen  in  the  air  acts  upon  the  dead 
products  of  animal  and  vegetable  life  and  converts  them 
into  harmless  substances.  In  this  way  it  acts  as  a  purify- 
ing agent.  For  example,  in  sewage-disposal  plants,  sewage 
is  forced  into  the  atmosphere  in  fine  sprays  (Fig.  13),  so 
that  the  oxygen  can  come  in  contact  with  the  putrid  matter 
in  the  sewage,  thus  purifying  the  sewage  and  preventing  it 


OXYGEN  23 

from  becoming  a  menace  to  health.  The  pure  commercial 
oxygen  is  also,  used  in  the  treatment  of  certain  diseases 
and  especially  as  a  source  of  intense  heat  (see  oxyhydrogen 
and  oxy acetylene  blowpipe). 

EXERCISES 

1.  In  Fig.  9,  why  does  the  water  stay  in  the  inverted  cylinder? 
Why  does  the  oxygen  displace  it  ?    When  a  little  oxygen  has  entered, 
why  does  not  all  the  water  run  out  ? 

2.  What  other  liquids  can  you  think  of  over  which  gases  might 
be  collected  ? 

3.  What  does  the  word  pneumatic  mean  (see  dictionary)  ? 

4.  Can  you  find  out  what  is  the  pressure  of  1  atmosphere  in 
pounds  per  square  inch?   in  grains  per  square  centimeter? 

5.  Since  oxygen  is  such  an  active  gas,  why  is  it  present  in  the 
atmosphere  in  such  large  quantities  ? 

6.  Can  combustion  take  place  without  the  emission  of  light? 

7.  Is  the  evolution  of  light  always  produced  by  combustion  ? 

8.  Why  are  oily  rags  more  likely  to  start  a  fire  than  oil  spilled 
on  the  floor  ? 

9.  From  the  percentages  given  (p.  15),  what  weight  of  potassium 
chlorate  will  be  required  to  yield  10  g.  of  oxygen?    Ans.  25.5  g. 

10.  What  weight  of  mercuric  oxide  will  be  required  to  yield  10  g. 
of  oxygen?    Ans.  135.13  g. 

11.  Assuming  the  cost  of  potassium  chlorate  to  be  50  cents  per 
kilogram  and  that  of  mercuric  oxide  to  be  $1.50  per  kilogram,  what 
is  the  cost  of  the  weight  of  each  required  in  the  preparation  of  1 0  g. 
of  oxygen?    Ans.    1.27  cents  and  20.27  cents. 

TOPICS  FOR  THEMES 

Joseph  Priestley  (Thorpe,  Essays  in  Historical  Chemistry). 
Spontaneous  combustion  (see  encyclopedia). 


CHAPTER  IV 
HYDROGEN 

Introduction.  A  great  variety  of  materials  undergo  com- 
bustion, among  them  being  coal,  wood,  oils,  and  various 
gases.  One  of  these  gases,  hydrogen,  is  of  special  interest 
because  it  is  an  elementary  substance. 

Various  combustible  gases  have  been  known  from  early 
ages,  but  they  were  long  confused  with  each  other.  The 
gas  hydrogen  was  first  clearly  recognized  as  a  distinct  sub- 
stance by  the  English  investigator,  Cavendish,  in  1776  ;  he 
obtained  it  in  pure  condition  and  showed  it  to  be  different 
from  all  other  known  gases.  It  was  named  hydrogen  by 
Lavoisier,  the  word  meaning  "  producer  of  water." 

Occurrence.  Hydrogen  occurs  in  the  atmosphere  in  the 
free  state,  but  only  in  traces.  It  occurs  in  enormous  quan- 
tities in  the  atmosphere  of  the  sun  and  certain  other  stars. 
In  the  combined  state  it  is  widely  distributed,  being  a  con- 
stituent of  water  as  well  as  of  all  living  organisms  and  of 
many  of  the  products  derived  from  them,  such  as  wood, 
starch,  and  sugar.  About  10  per  cent  of  the  human  body 
is  hydrogen.  Combined  with  carbon  it  forms  many  com- 
pounds, which,  mixed  together,  constitute  petroleum  and 
natural  gas. 

Preparation.  Hydrogen  can  be  prepared  in  a  number  of 
ways,  three  of  which  are  of  special  interest : 

1.  By  the  action  of  metals  on  water.  When  brought  into 
contact  with  certain  metals  under  appropriate  conditions, 
water  gives  up  the  whole  or  a  part  of  its  hydrogen,  its 

24 


HYDROGEN 


25 


place  being  taken  by  the  metal.  In  the  case  of  a  few  of 
the  metals  this  change  occurs  at  ordinary  temperatures. 
Thus,  if  a  bit  of  the  metal  sodium  is  dropped  on  water, 
an  action  is  seen  to  take  place  at  once,  sufficient  heat  being 
set  free  to  melt  the  sodium,  which  runs  about  on  the  sur- 
face of  the  water.  The  change  which  takes  place  consists 
in  the  substitution  of  one  half  of  the  hydrogen  of  the  water 
by  the  sodium,  and  may  be  represented  as  follows : 

sodium  +  water >-  hydrogen  +  sodium  hydroxide 

riiydrogenl  ["sodium     ~| 


|_oxygen     J 


hydrogen 
Loxygen     . 


The  sodium  hydroxide  formed  is  a  white  solid  which 
remains  dissolved  in  the  excess  of  undecomposed  water 
and  may  be  obtained  by  evaporating 
the  solution  to  dryness.  The  hydro- 
gen is  evolved  as  a  gas  and  may  be 
collected  by  suitable  means. 

Fig.  14  represents  a  simple  form  of  0=o 
apparatus  used  in  preparing  hydro- 
gen by  the  action  of  sodium  on  water. 
Since  the  sodium  is  lighter  than  water, 
it  is  kept  under  the  water  by  push- 
ing a  pellet  of  the  metal  into  the 
end  of  a  short  piece  of  lead  or  tin 
pipe,  the  other  end  of  which  has  been 
hammered  until  closed.  The  pipe  con- 
taining the  sodium  is  then  dropped 
into  a  trough  of  water.  Hydrogen 
is  at  once  evolved  and  is  collected  by  bringing  over  it  a  bottle 
or  cylinder  filled  with  water,  as  shown  in  the  figure. 

Other  metals,  such  as  magnesium  and  iron,  decompose 
water  rapidly  but  only  at  higher  temperatures.  When  steam 
is  passed  over  hot  iron,  for  example,  the  iron  combines  with 


FIG.  14.    The  preparation 

of  hydrogen  by  the  action 

of  sodium  on  water 


26 


FIRST  COUESE  IN  CHEMISTRY 


the  oxygen  of  the  steam,  setting  free  all  of  the  hydrogen. 
Experiments  show  that  the  change  may  be  represented  as 
follows : 

iron  -f-  water >•  hydrogen  -f-  iron  oxide 

Phydrogenl  Piron       "1 

Loxygen     J  LoxygenJ 

The  iron  oxide  formed  is  a  reddish-black  compound  identi- 
cal with  that  obtained  by  the  combustion  of  iron  in  oxygen. 


O 


FIG.  15.    The  preparation  of  hydrogen  by  the  action  of  iron  on  steam 

Preparation  of  hydrogen  from  iron  and  steam.  The  apparatus 
used  in  the  preparation  of  hydrogen  from  iron  and  steam  is 
shown  in  Fig.  15.  A  porcelain  or  iron  tube  A,  about  50  cm.  in 
length  and  2  cm.  or  3  cm.  in  diameter,  is  partly  filled  with  fine 
iron  wire  or  tacks  and  connected  as  shown  in  the  figure.  The 
tube  is  heated  slowly  at  first,  until  the  iron  is  red-hot.  Steam 
is  then  conducted  through  the  tube  by  boiling  the  water  in  the 
flask  B.  The  hot  iron  combines  with  the  oxygen  in  the  steam, 
setting  free  the  hydrogen,  which  is  collected  over  water  in  C. 

2.  By  the  action  of  metals  on  acids  (usual  laboratory 
method).  In  the  laboratory,  hydrogen  is  usually  prepared 
from  compounds  known  as  acids,  all  of  which  contain 


HYDROGEX 


27 


hydrogen.  When  acids  are  brought  in  contact  with  a 
number  of  the  different  metals,  the  latter  dissolve  and 
set  free  the  hydrogen  of  the  acid.  It  has  been  found  most 
convenient  and  economical,  in  preparing  hydrogen  by  this 
method,  to  use  either  zinc  or  iron  as  the  metal  and  either 
hydrochloric  or  sulfuric  acid  as  the  acid.  Hydrochloric  acid 
is  an  aqueous  solution  of  a  gaseous  compound  known  as 
hydrogen  chloride  (which  consists  of  2. 77  per  cent  hydrogen 
and  97.23  per  cent  chlorine),  while  sulfuric  acid  is  an  aque- 
ous solution  of  an  oily  liquid  known  as  hydrogen  sulfate 
(which  consists  of  2.05  per  cent  hydrogen,  32.70  per  cent 
sulfur,  and  65.25  per  cent  oxygen). 

The  changes  taking  place  in  the  preparation  of  hydrogen 
from  zinc  and  sulfuric  acid  may  be  represented  as  follows : 

zinc  4-  sulfuric  acid >-  zinc  sulfate  +  hydrogen 

[hydrogen"!  Pzinc       ~| 

sulfur  I  sulfur     I 

oxygen     J  LoxygenJ 

In  other  words,  the  zinc  takes  the  place  of  the  hydrogen 
in  sulfuric  acid.  The  resulting  compound 
contains  zinc,  sulfur,  and  oxygen  and  is 
known  as  zinc  sulfate.  This  remains  dis- 
solved in  the  water  in  the  acid.  It  may  be 

obtained  in  the 
form  of  a  white 
solid  by  evaporat- 
ing the  liquid  left 
after  the  metal 
lias  passed  into 
solution. 


FIG.  16.    The  preparation   of   hydrogen  by  the 
action  of  metals  on  acids 


Directions  for  preparing  hydrogen  from  acids.  The  preparation 
of  hydrogen  from  acids  is  carried  out  in  the  laboratory  as 
follows :  The  metal  is  placed  in  a  flask  or  wide-mouthed  bottle 


28  FIRST  COURSE  IN  CHEMISTRY 

A  (Fig.  16),  and  the  acid  is  added  slowly  through  the  funnel 
tube  B.  The  metal  dissolves  in  the  acid,  while  the  hydrogen 
which  is  liberated  escapes  through  the  exit  tube  C  and  is 
collected  over  water.  Pure  sulfuric  acid  will  not  act  readily 
upon  pure  zinc.  The  reaction  may  be  started,  however,  by  the 
addition  of  a  few  drops  of  a  solution  of  copper  sulfate. 

3.  Commercial  preparation.  In  preparing  hydrogen  on  a 
large  scale  for  commercial  uses  the  method  of  electrolysis 
of  acidulated  water  is  used,  as  explained  on  page  17. 

Properties  of  hydrogen.  Hydrogen  resembles  oxygen  in 
that  it  is  a  colorless,  tasteless,  odorless  gas.  It  is  the  light- 
est of  all  known  substances,  1  liter  of  the  gas  weighing 
only  0.08987  g.  Soap  bubbles  blown  with  hydrogen  rapidly 
rise  in  the  air. 

Hydrogen  is  more  difficult  to  liquefy  than  any  other 
gas,  with  the  exception  of  the  rare  gas  helium.  The  Eng- 
lish chemist,  Dewar,  however,  in  1898  succeeded  not  only 
in  obtaining  hydrogen  in  liquid  state  but  also  as  a  solid. 
Liquid  hydrogen  is  colorless.  Its  density,  that  is,  the  weight 
of  1  cc.,  is  only  0.07.  Its  boiling  point  is  —  252.7°  and  its 
melting  point  —  259°.  The  solubility  of  hydrogen  in  water 
is  very  slight,  being  still  less  than  that  of  oxygen. 

Pure  hydrogen  produces  no  injurious  results  when  in- 
haled. Of  course  one  could  not  live  in  an  atmosphere  of 
the  gas,  since  oxygen  is  essential  to  respiration. 

Chemical  conduct.  At  ordinary  temperatures  hydrogen 
is  not  an  active  element.  Under  suitable  conditions,  how- 
ever, it  combines  with  many  of  the  elements,  forming  com- 
pounds known  as  hydrides.  Thus,  hydrogen  and  chlorine, 
when  mixed  together,  will  combine  with  explosive  violence 
if  heated  or  if  exposed  to  the  sunlight.  The  product  formed 
in  either  case  is  called  hydrogen  chloride.  Under  suit- 
able conditions  hydrogen  combines  with  nitrogen  to  form 


HYDROGEN 


29 


ammonia,  and  with  sulfur  to  form  the  foul-smelling  gas 
hydrogen  sulfide.  At  ordinary  temperatures  hydrogen  and 
oxygen  may  be  mixed  without  action.  If  the  mixture  is 
heated  to  about  800°,  or  if  a  flame  is  brought  in  contact 
with  it,  a  violent  explosion  takes  place.  Nevertheless, 
under  proper  conditions  hydrogen  may  be  made  to  burn 
quietly  in  either  oxygen  or  air.  The  resulting  hydrogen 


FIG.  17.   Burning  hydrogen  and  collecting  the  product  of  its  combustion 

flame  is  almost  colorless  and  is  very  hot.  The  combus- 
tion of  the  hydrogen  is  due  to  its  union  with  oxygen  and 
the  product  of  the  combustion  is  an  oxide  of  hydrogen. 
That  this  compound  is  water  may  be  easily  shown  by 
experiment. 

Directions  for  burning  hydrogen.  The  combustion  of  hydrogen 
in  air  may  be  carried  out  safely  as  follows :  The  hydrogen 
is  generated  in  the  bottle  A  (Fig.  17),  is  dried  by  conducting 
it  through  the  tube  B  filled  with  some  substance  (usually  cal- 
cium chloride)  which  lias  a  great  attraction  for  moisture,  and 
escapes  through  the  tube  C,  the  end  of  which  is  drawn  out  to 
a  jet.  When  all  tin'  an-  has  been  expelled  from  the  apparatus, 
the  hydrogen  may  be  ignited.  It  then  burns  quietly,  since  only 


30 


FIRST  COURSE  IN  CHEMISTRY 


the  small  amount  of  it  which  escapes  from  the  jet  can  come  in 
contact  with  the  oxygen  of  the  air  at  any  one  time.  By  holding  a 
cold,  dry  bell  jar  or  bottle  over  the  flame  in  the  manner  shown  in 
the  figure  the  steam  formed  by  the  combustion  of  the  hydrogen 
is  condensed,  water  collecting  in  drops  on  the  sides  of  the  jar. 

Hydrogen  does  not  support  combustion.  While  hydrogen 
is  readily  combustible,  it  is  not  a  supporter  of  combustion ; 
in  other  words,  substances  will  not  burn  iu  it.  This  may 
be  shown  by  bringing  a  lighted  caudle  sup- 
ported by  a  stiff  wire  into  a  bottle  or  cylinder 
of  the  pure  gas,  as  shown  in  Fig.  18.  The 
hydrogen  is  ignited  by  the  flame  of  the  can- 
dle and  burns  at  the  mouth  of  the  cylinder, 
where  it  comes  in  contact  with  the  oxygen  in 
the  air.  When  the  candle  is  thrust  up  into 
the  gas,  its  flame  is  extinguished.  If  slowly 
withdrawn,  the  candle  is  relighted  as  it  passes 
through  the  layer  of  burning  hydrogen. 

Reduction.  On  account  of  its  tendency 
to  combine  with  oxygen,  hydrogen  has  the 
power  of  abstracting  it  from  many  of  its 
compounds.  Thus,  if  a  stream  of  hydrogen 
generated  in  A  (Fig.  19)  and  dried  by  pass- 
ing through  the  tube  B  (filled  with  calcium 
chloride)  is  conducted  through  the  tube  6',  which  contains 
some  copper  oxide  heated  to  a  moderate  temperature,  the 
hydrogen  abstracts  the  oxygen  from  the  copper  oxide.  The 
change  may  be  represented  as  follows : 


FIG.  18.    Hydro- 
gen extinguishes 
the    flame    of   a 
candle 


copper  oxide  +  hydrogen 

Tcopper  ~] 

L< 


.oxygen J 


->•  water  +  copper 

rhydrogen"! 
Loxygen     J 


The  water  formed  collects  in  the  cold  portions  of  the  tube 
C  near  its  end.    In  this  experiment  the  copper  oxide  is  said 


HYDEOGEN 


31 


to  undergo  reduction.  Reduction  may  therefore  be  defined 
as  the  process  of  withdrawing  oxygen  from  a  compound. 
As  we  shall  see,  the  term  reduction  is  also  used  with  a 
somewhat  different  meaning. 

Relation  of  reduction  to  oxidation.  At  the  same  time  that 
the  copper  oxide  is  reduced,  it  is  clear  that  the  hydrogen 
is  oxidized,  for  it  combines  wTith  the  oxygen  given  up  by 
the  copper  oxide. 
The  two  proc- 
esses are  there- 
fore very  closely 
related,  and  it 
usually  happens 
that  when  one 
substance  is  oxi- 
dized, some  other 
substance  is  re- 
duced. The  one 
which  gives  up 
its  oxygen  is  called  an  oxidizing  agent,  while  the  other,  which 
unites  with  the  oxygen  of  the  oxidizing  agent,  is  called  a 
reducing  agent. 

Uses  of  hydrogen.  Hydrogen  is  sometimes  used  as  a 
material  for  the  inflation  of  balloons,  but  usually  the  much 
cheaper  coal  gas  is  substituted  for  it,  and  even  hot  air  is 
used  when  the  duration  of  ascension  is  very  short.  It  has 
been  used  in  the  oxyhydrogen  blowpipe  as  a  source  of  light 
and  heat.  Where  the  electric  current  is  available,  however, 
the  blowpipe  has  been  displaced  almost  entirely  by  the  elec- 
tric light  and  electric  furnace,  which  are  more  economical 
and  more  powerful  sources  of  light  and  heat.  Its  greatest 
commercial  use  is  in  the  conversion  of  liquid  oils  into  solid 
edible  fats,  as  explained  later  (p.  239). 


FIG.  19.    Reduction  of  copper  oxide  by  hydrogen 


32  FIRST  COURSE  IN  CHEMISTRY 

The  oxy hydrogen  blowpipe.  This  is  a  form  of  apparatus  used 
for  burning  hydrogen  in  pure  oxygen.  It  consists  of  a  small 
tube  placed  within  a  larger  one,  as  shown  in  Fig.  20.  The 
hydrogen  is  first  passed  through  the  outer  tube  and  ignited  at 
the  open  end  of  the  tube  A.  The  oxygen  is  then  conducted 
through  the  inner  tube  and  mixes  with  the  hydrogen  at  the 

end  of  the  tube.  The  inten- 
sity of  the  heat  may  be 
shown  by  bringing  into  the 
flame  pieces  of  metal  such 
as  iron  wire  or  zinc.  These 
burn  with  great  brilliancy. 
Even  platinum,  having  a 
melting  point  of  1755°,  may 
be  melted  by  the  heat  of 
the  flame.  While  the  oxyhy- 
FIG.  20.  The  oxy  hydrogen  blowpipe  drogen  flame  is  intensely 

hot,   it  is   almost   nonlumi- 

nous.  If  it  is  directed  against  some  infusible  substance  like 
ordinary  lime,  the  heat  is  so  intense  that  the  lime  becomes  incan- 
descent and  glows  with  a  brilliant  light.  This  is  sometimes 
used  as  a  source  of  light  under  the  name  limelight. 


EXERCISES 

1.  In  Fig.  17  why  is  it  necessary  to  dry  the  hydrogen  by  means 
of  the  calcium  chloride  in  the  tube  Bl 

2.  From  Fig.  17  suggest  a  way  for  determining  experimentally 
the  quantity  of  water  formed  in  the  reaction. 

3.  In  Fig.  18  will  the  flame  remain  at  the  mouth  of  the  tube? 

4.  How  many  grams  of  hydrogen  can  be  made  from  100  g.  of 
hydrogen  sulfate  (p.  27)?    Am.  2.05  g. 

5.  How  many  grams  of  hydrogen  can  be  made  from  100  g.  of 
water?   Ans.  11.18g. 

TOPICS  FOR  THEMES 

Cavendish  (Thorpe,  Essays  in  Historical  Chemistry). 
Oxy  hydrogen  blowpipe  (see  encyclopedia). 


CHAPTER  V 


THE  GAS  LAWS;    STANDARD  CONDITIONS 

Introduction.  It  will  be  remembered  that  in  discussing 
the  properties  of  oxygen  and  hydrogen  the  weight  of  a  liter 
of  each  gas  was  given.  A  moment's  reflection  will  make 
it  clear  that  these  weights 
must  refer  to  some  set  of 
arbitrary  conditions,  for  it  is 
a  familiar  fact  that  the  vol- 
ume of  a  given  quantity  of  a 
gas  varies  both  with  changes 
in  pressure  and  with  changes 
in  temperature. 

Variation  of  volume  with 
pressure :  law  of  Boyle.  That 
the  volume  occupied  by  a 
given  weight  of  a  gas  can 
be  altered  by  changing  the 
pressure  is  familiar  to  every- 
one who  has  pumped  air  into 
a  bicycle  or  automobile  tire. 
As  early  as  1660  Robert 
Boyle,  an  Irish  investigator 

(Fig.  21),  reached  the  following  conclusion,  known  as 
Boyle's  law :  If  the  temperature  remains  constant,  the  volume 
occupied  by  a  given  weight  of  a  gas  is  inversely  proportional 
to  the  pressure.  Thus,  if  a  given  weight  of  a  gas  occupies 
a  volume  of  1000  cc.  when  subjected  to  a  certain  pressure, 

33 


FIG.  21.  Kobert  Boyle  (1626-1691) 

One  of  the  most  accurate  of  the  early 
chemists 


34 


FIKST  COURSE  IX  CHEMISTRY 


CENTIGRADE  AND  ABSO- 
LUTE SCALES 


it  will  occupy  a  volume  of  500  cc.  if  the  pressure  is 
doubled,  or  of  2000  cc.  if  the  pressure  is  diminished  to 
one  half.  This  means  that  for  a  given 
weight  of  a  gas  the  product  of  the 
pressure  into  the  volume  will  remain 
constant,  no  matter  how  either  one 
may  be  altered.  Designating  the  pres- 
sure and  volume  under  one  set  of 
conditions  by  P  and  F,  and  under  a 
different  set  by  1\  and  V^  Boyle's  law 
may  be  stated  thus  : 


Variation  of  volume  with  tempera- 
ture. If  the  pressure  is  held  constant, 
all  gases  expand  when  the  tempera- 
ture is  raised  and  contract  when  it  is 
lowered,  and  it  is  a  remarkable  fact 
that  the  volumes  of  all  gases  change 
to  the  same  extent  for  a  given  varia- 
tion in  the  temperature.  Let  us  sup- 
pose that  the  volume  of  a  gas  has  been 
measured  at  zero  on  the  centigrade 
scale.  Experiment  has  shown  that  a 
rise  of  one  degree  causes  an  expansion 
of  ^-i~g.  of  this  volume  ;  a  rise  of  five 
degrees,  an  expansion  of  -%§•$.  If  we 
take  273  cc.  of  this  gas  at  zero,  the 
volume  at  1°  above  will  be  274  cc.  ;  at 
1°  below  it  will  be  272  cc.  ;  and  at  5° 


Tin 

melts 

231.  9°  C. 

504.  9°  A. 

Water 

boils 

100°  C. 

373°  A. 

Phosphorus 

melts 

44°  C. 

317°  A. 

Water 

freezes 

0°C. 

273°  A. 

Mercury 

freezes 

-38.7°C. 

234.3°  A. 

Chlorine 

freezes 

-  101.5°  C. 

171.5°  A. 

Oxygen 

boils 

-  182.9°  C. 

90.1°  A. 

Absolute 

zero 

—  273°  C.     0°  A. 

FIG.  22.      Comparison 
of  temperatures  on  the 
centigrade    and    abso- 
lute scales 


below  it  will  be  268  cc.  At  the  same 
rate  of  contraction  the  volume  will  be  1  cc.  at  —  272°,  and 
at  —273°  it  will  be  zero.  Of  course  this  cannot  really 


THE  GAS  LAWS;  STAKDAKD  CONDITIONS     35 


happen,  and  experiment  shows  that  before  this  tempera- 
ture is  reached,  all  gases  have  changed  into  liquids  or 
solids.  Helium,  the  most  difficult  gas  to  liquefy,  passes 
into  a  liquid  at  -208.7°. 

The  absolute  scale  of  temperature.  If  we  were  to  con- 
struct a  thermometer  having  divisions  of  the  same  size  as 
those  on  the  centigrade  scale,  but  with  the  zero  point  at 
-  273°  on  the  latter  scale, 
then  the  point  at  which 
water  freezes  (0°  centigrade) 
would  be  273°.  At  272° 
on  this  scale  the  273  cc.  of 
gas  mentioned  in  the  last 
paragraph  would  measure 
272  cc. ;  at  271°  it  would 
measure  271  cc. ;  at  1°,  1  cc. 
On  such  a  scale  the  volume 
of  a  gas  would  be  propor- 
tional to  the  temperature  at 
every  point.  This  scale  is 
known  as  the  scale  of  abso- 
lute temperature,  the  point 
—  273°  centigrade  being  the 
absolute-zero  point.  Evidently 
the  absolute  temperature  may 
be  obtained  by  adding  273 
Fig.  22  gives  a  comparison  of  the  centigrade  and  absolute 
scales  at  a  number  of  temperatures. 

The  law  of  Gay-Lussac  (or  of  Charles).  A  general  state- 
ment can  now  be  made  in  regard  to  the  effect  of  temperature 
on  the  volume  of  a  gas :  If  the  pressure  remains  constant, 
the  volumes  occupied  ly  a  given  weight  of  a  gas  at  different 
temperatures  are  proportional  to  the  absolute  temperatures. 


FIG.  23.   Joseph  Louis  Gay-Lussac 
(1778-1850) 

A  distinguished  French  chemist 

to  the   centigrade  reading. 


36 


FIEST  COURSE  IN  CHEMISTRY 


If   V  and 
TV  then 


or 


PI  are  the  volumes  at  the  temperatures  T  and 
V  T 

T7__L1± 


The  above  generalization  is  called  the  law  of  Gay-Lussac 
(Fig.  23)  or  of  Charles,  since  it  was  formulated  independ- 
ently by  these  two  Frenchmen  in  1801. 
Standard  conditions.  It  is  now  easy 
to  choose  conditions  which  can  be 
regarded  as  standard  for  the  measure- 
ment of  gases,  and  these  will  be  taken 
for  granted  unless  a  special  statement 
to  the  contrary  is  made. 

As  a  standard  temperature  the  tem- 
perature of  melting  ice  is  chosen,  which 
is  0°  centigrade  or  273°  absolute. 

As  a  standard  pressure  the  average 
pressure  exerted  by  the  atmosphere  at 
sea  level  has  been  selected.  This  is 
equal  to  1033  g.  per  square  centi- 
meter. As  a  rule,  however,  the  pres- 
sure of  the  atmosphere  is  expressed 
by  stating  the  height  of  a  column  of 
mercury  which  the  pressure  of  the  atmosphere  will  sustain  ; 
and,  expressed  in  this  way,  the  standard  pressure  is  equal 
to  that  exerted  by  a  column  of  mercury  760  mm.  in  height. 
Aqueous  tension.  As  a  rule,  gases  are  measured  in  the 
laboratory  by  collecting  them  over  water  in  a  graduated 
tube,  as  represented  in  Fig.  24.  Before  the  reading  is  taken, 
the  tube  A  is  first  raised  or  lowered  until  the  level  of  the 
water  is  the  same  within  and  without  the  tube  ;  the  inclosed 
gas  is  then  under  atmospheric  pressure.  But  to  some  extent 


FIG.  24.   Measuring 

a  gas  collected  over 
water 


THE  GAS  LAWS;  STANDARD  CONDITIONS     37 

water  has  evaporated  into  the  tube,  and  a  part  of  the  vol- 
ume inclosed  is  due  to  water  vapor  and  not  to  the  gas.  If 
the  water  vapor  could  be  removed,  the  gas  would  occupy 
a  smaller  volume.  If  we  could  find  out  how  much  pressure 
the  water  vapor  exerts  upon  the  surface  of  the  water  within 
the  tube  and  subtract  this  from  the  atmospheric  pressure, 
we  should  have  the  pressure  which  the  gas  itself  exerts  at 
the  volume  which  it  occupies.  The  pressure  due  to  the 
water  vapor  is  called  the  aqueous  tension  or  the  vapor  pres- 
sure. It  increases  steadily  as  the  temperature  rises,  and  the 
table  in  the  Appendix  gives  its  value  over  the  range  of 
temperatures  most  required. 

Kinetic  molecular  theory.  It  is  a  surprising  fact  that  all 
gases  should  exert  a  considerable  pressure  when  confined 
in  a  closed  vessel,  and  that  the  volume  of  all  gases  should 
be  equally  affected  by  changes  in  pressure  and  temperature. 
These  facts  have  led  to  the  view  that  all  gases  are  made  up 
of  minute  particles,  called  molecules,  which  are  in  very  rapid 
motion.  Their  momentum,  as  they  strike  the  walls  of  the 
confining  vessel,  occasions  the  pressure  which  the  gas  exerts. 
This  conception  is  called  the  kinetic  molecular  theory. 

Examples  of  calculation.  A  few  typical  examples  of  the  appli- 
cation of  the  gas  laws  to  actual  calculations  will  make  their 
meaning  clear.  (It  will  be  understood  that  temperatures  are 
always  given  on  the  centigrade  scale  unless  otherwise  specified.) 

1.  A  gas  measured  under  a  pressure  of  720  mm.  had  a  volume 
of  620  cc.   What  volume  will  this  gas  occupy  under  standard 
pressure,  760  mm.,  the  temperature  remaining  constant  ? 

According  to  Boyle's  law,  PV=PlVr  Substituting  the 
values  given  in  the  problem,'  we  have  760  V  =  720  x  620 ;  or 
V  =  587.4  cc. 

2.  The  volume  of  a  gas  measured  at  a  temperature  of  90°  is 
930  cc.    What  will  be  its  volume  at  0°,  the  pressure  remaining 
constant  ? 


38  FIRST  COURSE  IN  CHEMISTRY 

According  to  the  law  of  Gay-Lussac,  V  :  F1 : :  T :  Tv  in  which 
T  and  1\  refer  to  the  absolute  scale.  Changing  the  centigrade 
temperatures,  given  in  the  problem,  to  absolute  temperatures 
by  adding  273,  and  substituting  in  the  formula,  we  have 

V :  930  : :  273°  :  363° ;  or  V  =  699.1  cc. 

3.  A  gas  measured  300  cc.  under  a  pressure  of  740  mm.  and 
a  temperature  of  25°  (or  298°  absolute  scale).  What  will  its 
volume  be  under  standard  conditions  (0°  and  760mm.  pressure)  ? 

First  find  the  change  in  volume  due  to  change  in  pressure : 

300  x  740  =  760  x  F;  or  F  =  292cc. 
Next  make  the  correction  for  temperature : 

292  :  F :  :  298  :  273 ;  or  F  =  264.5  cc. 

This  gives  the  volume  under  standard  conditions. 

When  the  gas  volumes  are  measured  over  water,  the  value 
of  the  aqueous  tension  (see  Appendix)  must  be  subtracted  from 
the  barometric  pressure.  For  example,  suppose  that,  in  prob- 
lem 3  above,  the  gas  had  a  volume  of  300  cc.  when  measured 
over  water  (Fig.  24) ;  the  pressure  would  not  then  be  740  mm. 
but  740  less  23.69,  which  is  the  value  for  the  aqueous  tension  at 
25°.  The -real  pressure,  therefore,  is  740-23.69,  or  716.31  mm. 
The  problem  is  then  solved  just  as  above,  except  that  the  value 
716.31  is  substituted  for  the  value  740. 

Practical  suggestions.  In  solving  such  problems  the  student 
should  notice  that  an  increase  of  pressure  diminishes  the  vol- 
ume, as  does  also  a  decrease  in  temperature.  After  solving  a 
problem,  he  should  compare  his  results  with  the  original  values 
and  see  if  they  are  reasonable.  It  will  be  of  assistance  to  re- 
member that  if  the  conditions  under  which  a  gas  is  measured  are 
actual  laboratory  conditions  (pressure  from  740  to  760  mm.  and 
temperature  from  15°  to  25°),  the  gas  will  contract  at  standard 
conditions  to  from  about  7  to  12  per  cent. 


THE  GAS  LAWS;  STANDARD  CONDITIONS      39 

EXERCISES 

1.  Give    two    illustrations    of     Boyle's    law    from    everyday 
experience. 

2.  Why  is  the  bottom  of  a  balloon  left  open  and  not  tightly 
closed  ? 

3.  How  does  the  air  brake  on  a  railway  car  work? 

4.  How  do  you  change  readings  on  the  Fahrenheit  scale  into 
centigrade  readings  ? 

5.  On  warm,  humid  days  an  automobile  engine  will  not  work 
as  smoothly  (without  adjustment)  as  on  cold,  dry  ones.    Why  is  this? 

6.  AVhy  does  a  balloon  tend  to  fall  toward  evening  and  rise  at 
midday  ? 

7.  A  gas  under  standard  pressure  measured  780  cc.  Temperature 
remaining  constant,  what  will   be  the  pressure  when  the  volume 
measures  360  cc.?    A ns.  1646.6  mm. 

8.  A  gas  at  0°  measured  560  cc.   Under  the  same  pressure  what 
will  be  its  volume  at  100°?   Am.  765.1  cc. 

9.  Under  standard  conditions  a  gas  measured  950  cc.   What  will 
be  its  volume  at  740  mm.  and  22°?    A  ns.  1054.3  cc. 

10.  A  gas  standing  over  water  at  20°  with  the  barometer  reading 
755  mm.  measured  100  liters.    Without  change  in  temperature  or 
barometric  pressure,  the  gas  was  passed  through  a  drying  agent  and 
collected  over  mercury.    What  was  its  volume?    Ans.  97.68  liters. 

11.  100  g.  of  potassium  chlorate  and  25  g.  of  manganese  dioxide 
were  heated  in  the  preparation  of  oxygen.    What  products  were  left 
in  the  flask,  and  how  much  of  each  was  present?  Ans.  60.8  g.  of 
potassium  chloride  and  25  g.  of  manganese  dioxide. 

TOPICS  FOR  THEMKS 

Robert  Boyle  (Thorpe,  Essays  in  Historical  Chemistry). 

Gay-Lussac  (see  encyclopedia). 

The  different  thermometers  in  common"  use  (see  encyclopedia). 


CHAPTER  VI 

COMPOUNDS  OF  HYDROGEN  AND  OXYGEN:    WATER  AND 
HYDROGEN  PEROXIDE 

WATER 

Historical.  Water  was  regarded  as  an  element  until 
1781,  when  Cavendish  showed  that  it  is  formed  by  the 
union  of  hydrogen  and  oxygen.  Being  a  believer  in  the 
phlogiston  theory,  however,  he  failed  to  interpret  his  results 
correctly.  A  few  years  later  Lavoisier  repeated  Cavendish's 
experiments  and  showed  that  water  must  be  regarded  as  a 
compound  of  hydrogen  and  oxygen. 

Occurrence  of  water.  Water  not  only  covers  about  five 
sevenths  of  the  surface  of  the  earth  and  is  present  in  the 
atmosphere  in  the  form  of  vapor,  but  it  is  also  a  common 
constituent  of  the  soil,  of  many  rocks,  and  of  almost  every 
form  of  animal  and  vegetable  organism.  Nearly  70  per  cent 
of  the  human  body  is  water.  This  is  derived  not  only  from 
the  water  which  we  drink  but  also  from  the  food  which  we 
eat,  most  of  which  contains  a  large  percentage  of  water.  The 
table  on  page  241  shows  the  percentage  of  water  present  in 
some  of  the  more  common  foods. 

Composition  of  natural  waters.  Water  as  it  occurs  in 
nature  always  contains  more  or  less  matter  derived  from 
the  rocks  and  soils  with  which  it  comes  in  contact.  When 
such  water  is  evaporated,  this  matter  is  left  behind  in 
solid  form.  Even  rain  water,  which  is  the  purest  natural 
water,  contains  dust  particles  and  gases  dissolved  from  the 

40 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN  41 

atmosphere.    The  foreign   matter   in   natural  waters   is  of 
two  kinds,  namely,  mineral  and  organic. 

1.  Mineral  matter.     The   mineral   substances   ordinarily 
present  in  fresh  waters  are  common  salt  and  compounds  of 
calcium,  magnesium,  and  iron.    Water  containing  any  con- 
siderable amounts  of  mineral  matter  does  not  form  a  lather 
with  soap,  and  is  termed  hard  water.  Water  containing  little 
or  no  mineral  matter,  such  as  rain  water,  is  termed  soft 
water.    One  liter  of  an  average  river  water  contains  about 
9.175  £.  of  mineral  matter.    The  waters  of  the  ocean  contain 

O 

about  40  g.  of  mineral  matter  to  the  liter,  more  than  three 
fourths  of  which  is  common  salt. 

2.  Organic  matter.    The  organic  matter  present  in  water 
consists  of  products  derived  from  animal  and  vegetable  life. 
Such  matter  is  absorbed  from  the  soil  or  introduced  from 
sewage.    Associated  with  such  matter  and  feeding  upon  it 
are  certain  living  microorganisms. 

Effect  of  the  foreign  matter  in  water  upon  health.  As 
a  rule,  any  sickness  resulting  from  drinking  impure  waters 
is  due  to  the  presence  of  living  microorganisms.  Many  of 
these  are  without  injurious  effect  upon  the  human  system, 
but  some  are  the  direct  cause  of  disease.  Thus,  a  transmis- 
sible disease  such  as  typhoid  fever  is  due  to  a  certain  kind 
of  organism  which,  through  food  or  drink,  is  introduced 
into  the  system.  It  is  easily  possible  for  these  organisms  to 
find  their  way,  through  sewage,  from  persons  afflicted  with  the 
disease  into  ivells  or  any  poorly  protected  water  supply,  and  it 
is  chiefly  in  this  ivay  that  typhoid  fever  is  spread. 

Purification  of  water.  Three  general  methods  are  used 
for  the  purification  of  water,  namely,  distillation,  boiling, 
and  filtration. 

1.  Distillation.  The  most  effective  way  of  purifying  ordi- 
nary water  is  by  the  process  of  distillation.  This  consists 


42 


FIRST  COURSE  IN  CHEMISTRY 


in  boiling  the  water  and  condensing  the  resulting  steam.    In 
the  laboratory  the  process  is  usually  conducted  as  follows : 

Ordinary  water  is  poured  into  the  flask  A  (Fig.  25)  and 
boiled.  The  steam  is  conducted  through  the  condenser  B,  com- 
monly known  as  a  Liebig  condenser,  which  consists  essentially 

of  a  narrow  glass 
tube  sealed  within 
a  larger  one,  the 
space  between  the 
two  being  filled 
with  cold  water, 
which  enters  at  C 
and  escapes  at  I). 
In  this  way  the 
inner  tube  is  kept 
cool  and  the  steam 
in  passing  through 
it  is  condensed. 
The  water  formed 
by  the  condensa- 
tion of  the  steam  collects  in  the  receiver  E  and  is  known  as 
distilled  water.  The  impurities  are  not  changed  into  vapor  but 
remain  in  the  flask  A. 

Distilled  water  is  pure  water.  It  is  used  by  the  chemist 
in  almost  all  of  his  work.  Large  quantities  are  also  used 
in  the  manufacture  of  ice,  as  well  as  for  drinking. 

Commercial  distillation.  In  preparing  distilled  water  on  a 
large  scale  the  steam  is  generated  in  a  metal  boiler  A  (Fig.  26) 
and  is  conducted  through  the  pipe  R  to  the  condensing  coil  C, 
made  of  tin.  This  pipe  is  wound  into  a  spiral  and  is  sur- 
rounded by  cold  water,  which  enters  at  D  and  flows  out  at  E. 
The  distilled  water  is  collected  in  a  suitable  container  F. 

2.  Boiling.  In  purifying  water  for  drinking  purposes  it 
is  only  necessary  to  remove  or  destroy  the  microorganisms 


FIG.  25.   The  distillation  of  water  in  the 
laboratory 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN  43 

present.  When  the  amount  of  water  to  be  purified  is  small, 
as  is  the  case  with  the  household  supply  for  drinking,  this 
is  most  conveniently  accomplished  by  boiling  the  water  for 
ten  or  fifteen  minutes.  While  the  organisms  are  destroyed 
in  a  short  time  by  moist  heat,  even  severe  cold  has  been 
found  to  have  comparatively  little  effect  upon  them. 


FIG.  26.   The  commercial  distillation  of  water 

3.  Filtration.  On  a  small  scale,  water  is  filtered  in  two 
ways :  (1)  by  passing  it  through  some  porous  material,  such 
as  charcoal,  or  (2)  by  forcing  it  through  porous  clay  ware, 
as  is  done  in  the  Chamberlain-Pasteur  filter.  While  such 
filters,  if  kept  clean  and  in  good  condition,  remove  most 
of  the  organic  matter,  they  do  not  remove  mineral  matter 
except  such  as  is  held  in  suspension.  These  household  filters 
are  not  easily  kept  in  order  and  soon  become  ineffective. 
They  are  no  longer  used  to  any  great  extent. 


44 


FIRST  COURSE  IN  CHEMISTRY 


City  filtration.  Many  cities  find  it  necessary  to  take  their 
water  supply  from  rivers.  The  rivers,  especially  in  thickly 
populated  districts,  are  almost  certain  to  be  contaminated  with 
organic  matter,  suggesting  the  possible  presence  of  disease 
germs.  Such  water  is  a  constant  menace  to  the  health  of  the 
city,  so  that  it  is  of  the  greatest  importance  to  find  some  way 
of  purifying  it  effectively  on  a  large  scale.  This  is  done  by 
filtration.  Two  general  kinds  of  filters  are  in  use : 

1.  Slow  sand  filters  (Fig.  27).  These  consist  of  large  beds  of 
sand  and  gravel,  through  which  the  water  passes  slowly.  Some 

of  the  impuri- 
ties are  strained 
out,  while  others 
are  decomposed 
by  the  action  of 
certain  kinds  of 


FIG.  27.   A  covered  sand-filter  bed 


which  collect  in 
a  jelly  like  layer 
on  the  surface 
of  the  filter.  The 
purified  water 
passes  into  the 
porous  pipe  A, 
from  which  it  is 
pumped  into  the 

city  mains.    The  filters  are  covered  to  protect  the  water  and 
prevent  it  from  freezing. 

2.  Mechanical  filters.  In  these  the  water,  before  filtration, 
is  run  into  large  tanks  and  treated  with  certain  compounds, 
such  as  aluminium  sulfate  or  iron  sulfate,  which  form  in  the 
water  a  small  amount  of  gelatinous  solid.  This  slowly  settles 
to  the  bottom  (p.  328),  carrying  with  it  much  of  the  organic 
matter  present.  The  partially  clarified  water  is  then  filtered 
through  sand  and  gravel.  In  the  United  States  about  400 
cities  and  towns  are  filtering  their  water  supply,  350  of  which 
are  using  the  mechanical  filters. 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN  45 


The  effect  of  the  filtration  of  the  water  supply  upon  the 
health  of  a  city  is  shown  by  the  fact  that  in  general  the  num- 
ber of  cases  of  typhoid  fever  in  cities  which  have  introduced 
an  effective  water-purification  system  has  been  decreased  by 
about  75  per  cent.  The  number  of  cases  of  many  other  diseases 
has  been  likewise  diminished. 

Self-purification  of  water.  It 
has  long  been  known  that 
water  contaminated  with  or- 
ganic matter  tends  to  purify 
itself  when  exposed  to  the 
air  (p.  22).  This  is  due  to 
the  fact  that  air  is  somewhat 
soluble  in  water  and  that  the 
dissolved  oxygen,  in  the  pres- 
ence of  certain  microorgan- 
isms, gradually  oxidizes  the 
organic  matter  present  in  the 
water  ;  when  this  is  destroyed, 
the  organisms  present  die  for 
lack  of  food.  While  water  is 
undoubtedly  purified  in  this 
way,  the  process  cannot  be 
relied  upon  to  purify  a  con- 
taminated water  so  as  to  render 
it  safe  for  drinking  purposes. 


FIG.  28.   Justus  Liebig  (180<M873) 

A  great  German  chemist  and  teacher. 

A  pioneer,  especially  in  agricultural 

chemistry 


Properties  of  water.  Pure 
water  is  an  odorless  and  taste- 
less liquid,  colorless  in  thin  layers  but  having  a  bluish 
tinge  when  observed  through  a  considerable  thickness.  It 
solidifies  at  0°  and  boils  at  100°  under  the  normal  pressure 
of  1  atmosphere.  When  water  is  cooled,  it  steadily  con- 
tracts until  the  temperature  of  4°  is  reached ;  at  lower  tem- 
peratures it  expands.  Water  is  remarkable  for  its  ability 
to  dissolve  other  substances,  and  is  the  most  general  solvent 


46  FIRST  COURSE  IN  CHEMISTRY 

known.  Chemists  usually  employ  aqueous  solutions  of  sub- 
stances rather  than  the  substances  themselves,  since  as  a 
rule  chemical  action  takes  place  more  readily  in  solution. 

Chemical  conduct.  Water  is  a  very  stable  substance ;  in 
other  words,  it  does  not  undergo  decomposition  readily. 
To  decompose  it  into  its  elements  by  heat  alone  requires 
a  very  high  temperature.  Even  at  2500°  only  about  10  per 
cent  of  the  water  heated  is  decomposed.  Though  very  stable 
toward  heat,  water  can  be  decomposed  in  other  ways,  as  by 
the  action  of  the  electric  current  or  by  certain  metals. 

Though  containing  88.81  per  cent  of  oxygen,  water  is 
not  a  good  oxidizing  agent,  because  of  its  great  stability. 
However,  certain  metals,  as  well  as  carbon,  can  be  oxidized 
by  very  hot  steam,  the  hydrogen  being  set  free.  Water  com- 
bines directly  with  many  compounds,  forming  substances 
called  hydrates.  Blue  vitriol  and  alum  are  good  examples 
of  such  hydrates. 

Heat  of  formation  and  heat  of  decomposition  are  equal.  The 
fact  that  a  very  high  temperature  is  necessary  to  decompose 
water  into  hydrogen  and  oxygen  is  in  accord  with  the  fact  that 
a  great  deal  of  heat  is  evolved  by  the  union  of  hydrogen  and 
oxygen  (p.  32),  for  it  has  been  proved  that  the  heat  necessary 
to  decompose  a  compound  into  its  elements  (heat  of  decom- 
position) is  equal  to  the  heat  evolved  in  the  formation  of  the 
same  compound  from  its  elements  (heat  of  formation). 

The  determination  of  the  exact  composition  of  water.  To 
determine  the  quantitative  composition  of  a  compound,  such 
as  water,  we  must  first  ascertain  what  particular  elements 
are  present  in  it  and  then  the  proportion  in  which  these 
elements  are  united.  We  have  already  shown  that  water  is 
composed  of  hydrogen  and  oxygen  (Figs.  17  and  19).  It 
remains  for  us  to  determine  in  what  ratio  these  elements 
are  combined  in  the  compound. 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN  4T 


The  proportion  in  which  hydrogen  and  oxygen  combine  to 
form  water.  By  mixing  known  volumes  of  hydrogen  and 
oxygen,  causing  them  to  combine,  and  then  ascertaining  the 
volume  and  identity  of  the  gas  remaining,  it  is  possible  to 
determine  the  exact  proportion  in  which  the  two  gases  com- 
bine to  form  water.  Such  a  process  is  called  a  synthesis,  and 
it  may  be  carried  out  as  follows : 

Details  of  the  experiment.  The  combination  of  the  two  gases 
is  brought  about  in  a  tube  called  a  eudiometer.  This  is  a 
graduated  glass  tube  about  60  cm.  long  and 
2  cm.  wide,  closed  at  one  end  (Fig.  29). 
Near  the  closed  end  two  platinum  wires 
are  fused  through  the  glass,  the  ends  of  the 
wires  within  the  tube  being  separated  by  a 
space  of  2  or  3  mm.  The  tube  is  entirely 
filled  with  mercury  and  inverted  in  a  ves- 
sel of  the  same  liquid.  Pure  hydrogen  is 
passed  into  the  tube  until  it  is  about  one 
fourth  filled.  The  tube  is  then  lowered  until 
the  mercury  stands  at  the  same  level  inside 
and  outside  the  tube,  and  the  reading  of  the 
volume  of  the  hydrogen  is  taken.  Approx- 
imately an  equal  volume  of  pure  oxygen 
is  then  introduced,  and  the  volume  is  again 
taken.  This  gives  the  total  volume  of  the 
two  gases.  From  this  the  volume  of  the  oxy- 
gen introduced  may  be  determined  by  sub- 
tracting from  it  the  volume  of  the  hydrogen. 

The  combination  of  the  two  gases  is  now  brought  about  by 
connecting  the  two  platinum  wires  with  an  induction  coil  and 
passing  a  spark  from  one  wire  to  the  other.  Immediately  a  slight 
explosion  occurs.  The  mercury  in  the  tube  is  at  first  depressed 
because  of  the  expansion  of  the  gases  due  to  the  heat  generated, 
but  it  at  once  rebounds,  taking  the  place  of  the  gases  which  have 
combined  to  form  water.  The  volume  of  the  water  in  the  liquid 
state  is  so  small  that  it  may  be  disregarded  in  the  calculations. 


FIG.  29.     The    eudi- 
ometer  employed  in 
determining  the  com- 
position of  water 


48  FIRST  COURSE  IN  CHEMISTRY 

In  order  that  the  temperature  of  the  residual  gas  and  the 
mercury  may  become  uniform,  the  apparatus  is  allowed  to  stand 
for  a  few  minutes,  and  the  volume  of  the  gas  is  taken.  The 
residual  gas  is  then  tested  in  order  to  ascertain  whether  it  is 
hydrogen  or  oxygen,  since  experiments  have  proved  that  it  is 
never  a  mixture  of  the  two.  From  the  information  thus  obtained 
the  composition  of  the  water  may  be  calculated. 

Calculation  of  composition.  Thus,  suppose  the  readings  were 
as  follows : 

Volume  of  hydrogen 20.3  cc. 

Volume  of  hydrogen  and  oxygen 38.7  cc. 

Volume  of  oxygen 18.4  cc. 

Volume  of  gas  left  after  combination  has  taken 

place  (found  to  be  oxygen) 8.3  cc. 

We  have  thus  found  that  20.3  cc.  of  hydrogen  have  com- 
bined with  18.4  cc.  minus  8.3  cc.  (or  10.1  cc.)  of  oxygen ;  or 
approximately  2  volumes  of  hydrogen  have  combined  with 
1  volume  of  oxygen.  Since  oxygen  is  15.9  times  as  heavy  as 
hydrogen,  the  proportion  by  weight  in  which  the  two  gases 
combine  is  1  part  of  hydrogen  to  7.94  parts  of  oxygen. 

Method  used  by  Berzelius  and  Dumas.  The  method  used 
by  these  investigators  enables  us  to  determine  directly  the 
proportion  by  weight  in  which  the  hydrogen  and  oxygen 
combine. 

Details  of  the  experiment.  Fig.  30  illustrates  the  essential 
parts  of  the  apparatus  used  in  making  the  determination.  The 
glass  tube  B  contains  copper  oxide,  while  the  tubes  C  and  D 
are  filled  with  calcium  chloride,  a  substance  which  has  great 
affinity  for  water.  The  tubes  B  and  C,  including  their  con- 
tents, are  carefully  weighed,  and  the  apparatus  is  connected  as 
shown  in  the  figure.  A  slow  current  of  pure  hydrogen  is  then 
passed  through  A,  and  that  part  of  the  tube  B  which  contains 
copper  oxide  is  carefully  heated.  The  hydrogen  combines  with 
the  oxygen  of  the  copper  oxide  to  form  water,  which  is  absorbed 
by  the  calcium  chloride  in  tube  C.  The  calcium  chloride  in 


COMPOUNDS  OF  HYDKOGEN  AND  OXYGEN  49 

tube  D  prevents  any  moisture  entering  tube  C  from  the  air. 
The  operation  is  continued  until  an  appreciable  amount  of 
water  has  been  formed.  The  tubes  B  and  C  are  then  weighed 
once  more.  The  loss  of  weight  in  the  tube  B  will  exactly  equal 
the  weight  of  oxygen  taken  up  from  the  copper  oxide  in  the 
formation  of  the  water.  The  gain  in  weight  in  the  tube  C 
will  exactly  equal  the  weight  of  the  water  formed.  The  differ- 
ence in  these  weights  will  of  course  equal  the  weight  of  the 
hydrogen  present  in  the  water  formed  during  the  experiment. 


FIG.  30.   Apparatus  employed  in  determining  the  ratio  by  weight  in  which 
oxygen  and  hydrogen  combine 


Dumas's  results.  The  above  method  for  the  determina- 
tion of  the  composition  of  water  was  first  used  by  Berzelius 
in  1820.  The  work  was  repeated  in  1843  by  Dumas,  who 
obtained  the  following  results : 


Weight  of  water  formed 

Oxygen  given  up  by  the  copper  oxide 
Weight  of  hydrogen  present  in  water 


945.439  g. 
840.161  g. 
105.278  g. 


According  to   this   experiment  the   ratio  of  hydrogen  to 
oxygen  in  water  is  105.278  :  840.161,  or  1 :  7.98. 

Morley's  results.  In  recent  years  the  American  chemist 
Morley  has  determined  the  composition  of  water  with 
great  care.  Extreme  precautions  were  taken  to  use  pure 
materials  and  to  eliminate  all  sources  of  error.  The  hydro- 
gen and  oxygen  which  combined,  as  well  as  the  water 


50 


FIRST  COUESE  IN  CHEMISTRY 


formed,  were  all  accurately  weighed.  According  to  Mor- 
ley's  results,  1  part  by  weight  of  hydrogen  combines  with 
7.94  parts  by  weight  of  oxygen  to  form  water. 

Comparison  of  results  obtained.  From  the  above  discus- 
sion it  is  easy  to  see  that  it  is  by  experiment  alone  that 
the  composition  of  a  compound  can  be  determined.  Differ- 
ent methods  may  lead  to  slightly  different  results.  The 
more  accurate  the  method  chosen,  and  the  greater  the  skill 
with  which  the  experiment  is  car- 
ried out,  the  more  accurate  will 
be  the  results.  It  is  generally  con- 
ceded by  chemists  that  the  results 
obtained  by  Morley  in  reference  to 
the  composition  of  water  are  the 
most  accurate  ones.  In  accordance 
with  these  results,  then,  water  must 
be  regarded  as  a  compound  con- 
taining hydrogen  and  oxygen  in 
the  ratio  of  1  part  by  weight  of 
hydrogen  to  7.94  parts  by  weight 
of  oxygen. 

Relation  between  the  volume  of 
aqueous  vapor  and  the  volumes  of 
the  hydrogen  and  oxygen  which  com- 
bine to  form  it.  If  the  quantitative  synthesis  of  water  as 
described  above  (Fig.  29)  is  carried  out  at  a  temperature 
above  100°,  the  water  vapor  formed  is  not  condensed,  and 
it  thus  becomes  possible  to  compare  the  volume  of  the 
water  vapor  with  the  volumes  of  hydrogen  and  oxygen 
which  combined  to  form  it.  This  can  be  accomplished  by 
surrounding  the  upper  part  of  the  eudiometer  A  (Fig.  31) 
with  a  glass  tube  J5,  through  which  is  passed  at  C  the  vapor 
obtained  by  boiling  some  liquid  which  has  a  boiling  point 


FIG.  31.  Eudiometer  for 
measuring  the  volume  of 
steam  formed  by  the  union 
of  oxygen  and  hydrogen 


COMPOUNDS  OF  HYDROGEN  AND  OXYGEN  51 

above  100°.  This  vapor  keeps  the  tube  A  heated  above  the 
boiling  point  of  water.  In  this  way  it  has  been  proved  that 
2  volumes  of  hydrogen  and  1  volume  of  oxygen  combine 
to  form  exactly  2  volumes  of  water  vapor.  It  will  be  noted 
that  the  relation  between  these  volumes  may  be  expressed  by 
whole  numbers.  The  significance  of  this  very  important  fact 
will  be  discussed  in  a  subsequent  chapter. 
^  Law  of  definite  composition.  We  have  just  seen  that 
water  contains  hydrogen  and  oxygen  combined  in  a  per- 
fectly definite  ratio.  In  the  earlier  days  of  chemistry  there 
was  much  discussion  as  to  whether  the  composition  of  a 
given  compound  is  always  precisely  the  same  or  whether 
it  is  subject  to  some  variation.  Experiments  have  shown, 
however,  that  the  composition  of  a  pure  chemical  com- 
pound is  always  exactly  the  same.  Thus,  pure  water 
obtained  from  any  source  whatever,  such  as  melting  pure 
ice,  condensing  steam,  or  burning  hydrogen  in  oxygen, 
always  contains  1  part  by  weight  of  hydrogen  to  7.94  parts 
of  oxygen.  This  truth  is  known  as  the  law  of  definite 
composition  and  may  be  stated  thus :  The  composition  of 
a  chemical  compound  never  varies. 

HYDROGEN  PEROXIDE 

Composition.  As  has  been  shown,  1  part  by  weight  of 
hydrogen  combines  with  7.94  parts  by  weight  of  oxygen 
to  form  water.  It  is  possible,  however,  to  obtain  a  second 
compound  of  hydrogen  and  oxygen  differing  from  water 
in  composition  in  that  1  part  by  weight  of  hydrogen  is 
combined  with  2  x  7.94,  or  15.88,  parts  of  oxygen.  This 
compound  is  called  hydrogen  peroxide,  the  prefix  per-  signi- 
fying that  it  contains  more  oxygen  than  hydrogen  oxide, 
which  is  the  chemical  name  for  water. 


52  FIRST  COURSE  IN  CHEMISTRY 

Preparation.  Hydrogen  peroxide  is  prepared  by  the  action 
of  acids  upon  barium  peroxide.  When  sulfuric  acid  is  used, 
the  change  which  takes  place  may  be  indicated  as  follows : 

barium    "]        f  sulfuric  1  { barium  1        f  hydrogen 

\  ^"   1  r   -\~  •( 

peroxide  J        ^  acid  [  sulfate  j        [  peroxide 

["hydrogen"! 

Loxygeu     _ 


f  barium  "I  Thydrogen"!  [barium  | 

oxygenj  sulfur  sulfur 

[oxygen  [oxygen J 


Properties  and  chemical  conduct.  Hydrogen  peroxide  is 
a  clear,  sirupy  liquid  having  a  density  of  1.458.  It  is  diffi- 
cult to  prepare  in  a  pure  state,  since  it  is  very  unstable,  de- 
composing into  water  and  oxygen  with  explosive  violence  : 

hydrogen  peroxide >•  water  -f-  oxygen 

In  dilute  solution  it  is  fairly  stable,  although  it  should  be 
kept  in  a  dark,  cool  place ;  otherwise  the  solution  loses 
its  strength,  the  hydrogen  peroxide  present  gradually  de- 
composing into  water  and  oxygen.  The  presence  of  a  small 
percentage  of  certain  substances,  such  as  a  trace  of  acid, 
preserves  the  strength  by  retarding  decomposition. 

Uses.  Solutions  of  hydrogen  peroxide  are  used  largely 
as  oxidizing  agents.  The  solution  sold  by  druggists  con- 
tains 97  per  cent  water  and  3  per  cent  of  the  peroxide, 
and  is  used  in  medicine  as  an  antiseptic.  Its  use  as  an 
antiseptic  depends  upon  its  oxidizing  properties.  It  acts 
upon  certain  dyes  and  natural  colors,  such  as  that  of  the 
hair,  oxidizing  them  to  colorless  compounds  ;  hence  it  is 
sometimes  used  as  a  bleaching  agent.  A  stronger  solution 
(30  per  cent)  is  used  as  an  oxidizing  agent  in  certain 
chemical  operations. 

The  law  of  multiple  proportion.  It  has  been  shown  that 
both  water  and  hydrogen  peroxide  are  compounds  of 
hydrogen  and  oxygen,  and  that  the  ratio  by  weight  in 


COMPOUNDS  OF  HYDKOGEN  AND  OXYGEN  53 

which  these  two  elements  are  present  in  each  of  these 
compounds  is  as  follows: 

Water hydrogen  :  oxygen  :  :  1  :  7.94 

Hydrogen  peroxide     .     .     .     hydrogen  :  oxygen  :  :  1 :  15.88 

It  will  be  seen  that  the  ratio  between  the  weights  of  oxy- 
gen combined  with  a  fixed  weight  of  hydrogen  (say  1  g.) 
in  these  two  compounds  is  7.94  :  15.88,  or  1  :  2. 

Similarly,  many  elements  other  than  oxygen  and  hy- 
drogen unite  to  form  a  number  of  distinct  compounds, 
each  with  its  own  precise  composition.  In  all  such  com- 
pounds the  same  statement  holds  as  in  the  case  of  water 
and  hydrogen  peroxide  —  the  weights  of  the  one  element 
which  are  combined  with  a  fixed  weight  of  the  other 
always  bear  a  simple  ratio  to  each  other,  such  as  1:2  or 
2  :  3.  This  truth  is  known  as  the  law  of  multiple  proportion. 
It  was  formulated  by  John  Dalton  (p.  67)  in  1808,  and 
may  be  stated  thus:  Wlwn  any  two  elements,  A  and  B, 
combine  to  form  more  than  one  compound,  the  weights  of 
A  which  unite  with  any  fixed  weight  of  B  bear  the  ratio  of 
small  whole  numbers  to  each  other. 


EXERCISES 

1.  In  making  solutions  why  does  the  chemist  use  distilled  water 
rather  than  filtered  water  ? 

2.  How  could  you  determine  the  total  amount  of  solid  matter 
dissolved  in  a  sample  of  water? 

3.  How  could  you  determine  whether  a  given  sample  of  water  is 
distilled  water  ? 

4.  How  could  the  presence  of  air  dissolved  in  water  be  detected? 

5.  How  could  the  amount  of  water  in  a  food  such  as  bread  or 
potato  be  determined? 

6.  Would  ice  frozen  from  impure  water  necessarily  be  free  from 
disease  germs  ? 


54  FIRST  COURSE  IN  CHEMISTRY 

7.  Suppose    that  the    maximum   density   of    water  were  at  0° 
instead  of  at  4°.    What  effect  would  this   have  on  the  formation 
of  ice  on  bodies  of  water? 

8.  W  hy  is  it  that  merely  heating  water  to  the  boiling  point  is  not 
sufficient  to  render  it  safe  for  sanitary  purposes  ? 

9.  If  steam  is  heated  to  2000°  and  again  cooled,  has  any  chemi- 
cal action  taken  place  ? 

10.  Why  is  cold  water  passed  into  C  instead  of  D  (Fig.  25)? 

11.  Mention  at  least  two  advantages  that  a  metal  condenser  has 
over  a  glass  condenser. 

12.  Draw  a  diagram  of  the  apparatus  used  in  your  laboratory  for 
supplying  distilled  water. 

13.  20  cc.  of  hydrogen  and  7  cc.  of  oxygen  are  placed  in  a  eudi- 
ometer and  the  mixture  exploded.    How  many  cubic  centimeters  of 
aqueous  vapor  are  formed?   Ans.  14  cc.   What  gas  and  how  much  of 
it  remains  in  excess  ?    A  ns.  6  cc.  hydrogen. 

14.  What  weight  of   oxygen  is  contained  in   100  g.  of   water? 
Ans.  88.81  g.     In  100  g.  of  pure  hydrogen  peroxide?    Ans.  94.07  g. 

TOPICS  FOR  THEMES 

Methods  used  in  your  city  for  obtaining  pure  water. 

Comparison  of  the  properties  of  water  with  those  of  the  gases 
from  which  it  is  formed. 


CHAPTER  VII 
MATTER  AND  ENERGY 

Definition  of  matter.  The  term  matter  includes  every- 
thing that  occupies  space  and  possesses  weight  or  mass. 
The  word  therefore  applies  equally  to  the  gases  of  the 
atmosphere,  the  water  of  the  ocean,  and  all  the  solid  objects 
which  make  up  the  world. 

Conservation  of  matter.  The  facts  we  have  learned  about 
burning,  and  the  experiments  we  have  made  with  oxygen 
and  hydrogen,  have  taught  us  that  matter  may  undergo 
many  transformations  as  a  result  of  chemical  action,  and 
that  many  of  its  properties  change  as  a  result  of  such  action. 
One  property  alone  never  changes,  and  that  is  the  mass. 
This  very  important  fact  was  first  clearly  recognized  by 
Lavoisier,  and  a  general  statement  of  it  is  known  as  the 
law  of  conservation  of  matter.  The  law  may  be  stated 
thus :  During  all  the  changes  through  which  a  given  quantity 
of  matter  may  pass,  its  mass  remains  constant. 

States  of  matter.  We  have  found  that  water  may  appear 
in  three  very  different  conditions,  depending  upon  the 
temperature,  namely,  solid,  liquid,  and  gaseous.  These  are 
called  the  three  states  of  matter.  This  is  not  a  peculiarity 
of  water,  for  all  substances  exist  in  all  three  states,  save 
only  when  the  temperature  required  for  the  melting  of  the 
solid  or  the  vaporization  of  the  liquid  is  so  high  that  de- 
composition takes  place  before  the  change  is  effected.  Thus, 
we  have  seen  that  mercuric  oxide  decomposes  before  it 

55 


56  FIKST  COURSE  IN  CHEMISTEY 

melts.   Potassium  chlorate  can  be  melted  without  difficulty, 
but  it  decomposes  before  it  boils. 

The  freezing  and  melting  points.  A  solid  normally  passes 
into  a  liquid  at  a  perfectly  definite  temperature,  called  its 
melting  point.  A  given  weight  of  any  solid,  in  melting,  ab- 
sorbs a  definite  quantity  of  heat,  the  exact  amount  absorbed 
depending  upon  the  solid.  This  is  known  as  the  heat  of  fusion. 
On  the  other  hand,  the  liquid  formed  tends  to  pass  back  into 
the  solid  state  at  this  same  temperature,  called  \\\Q  freezing 
point,  and  in  so  doing  the  heat  of  fusion  is  given  out  again. 
In  this  case  it  is  called  the  heat  of  solidification.  For  example, 
water  freezes  and  ice  melts  at  the  same  temperature,  namely, 
0°.  Moreover,  1  g.  of  ice  at  0°,  in  melting,  absorbs  a  quan- 
tity of  heat  that  would  raise  the  temperature  of  1  g.  of 
water  from  0°  to  79°,  while  1  g.  of  water  at  0°,  in  freezing, 
gives  out  this  same  quantity  of  heat. 

Sometimes  it  is  possible  to  cool  the  liquid  below  the  freezing 
point,  and  it  is  then  said  to  be  undercooled.  If  a  fragment  of 
the  solid  is  placed  in  the  undercooled  liquid,  solidification  will 
at  once  begin,  and  the  temperature  will  rise  to  the  true  freezing 
point  and  remain  there  as  solidification  continues.  The  freez- 
ing point  is  therefore  best  defined  as  the  temperature  at 
which  the  liquid  and  the  solid  can  be  mixed  without  change 
in  temperature. 

Vaporization.  There  is  no  definite  temperature  at  which 
a  liquid  passes  into  a  vapor,  or  gas.  Water  exposed  to  the 
air  vaporizes  at  all  temperatures,  and  even  ice  and  snow 
evaporate  during  weather  in  which  no  melting  occurs.  The 
higher  the  temperature  the  more  rapid  the  evaporation. 

Boiling  point.  The  escape  of  vapor  from  an  open  vessel 
containing  a  liquid  is  hindered  by  the  air  which  presses 
upon  the  surface  of  the  liquid.  The  vapor  escapes  by 
making  its  way  slowly  through  the  air,  but  when  the 


MATTER  AND  ENERGY 


57 


vapor  pressure  just  exceeds  the  pressure  exerted  by  the 
air,  there  is  nothing  to  prevent  the  vapor  from  making  its 
escape  as  fast  as  it  forms,  pushing  the  air  before  it.  When 
this  is  the  case,  any  additional  heat  applied  to  the  liquid 
will  not  raise  its  temperature,  but  will  merely  increase  its 
rate  of  evaporation. 

This  temperature  is  called  the  boiling  point  of  the  liquid, 
and  the  boiling  point  may  be  defined  as  the  temperature  at 
which  the  vapor  pressure  of  the  liquid  just  exceeds  the  pressure 
of  the  atmosphere.  It  will  be 
noticed  that  the  boiling  point 
of  a  liquid  is  not  fixed,  but 
depends  upon  the  atmospheric 
pressure. 

Liquefaction  of  gases.  Since 
increase  of  pressure  tends  to 
raise  the  boiling  point  of  a 
liquid,  it  is  clear  that  pres- 
sure applied  to  a  gas  will 
tend  to  condense  it  into  a 
liquid.  Faraday  (Fig.  33),  be- 
ginning about  1832,  was  the 
first  to  apply  this  principle  in  an  attempt  to  liquefy  gases, 
and  he  succeeded  in  liquefying  many  gaseous  substances 
which  up  to  that  time  had  never  been  prepared  in  the 
liquid  state. 

Method  of  Faraday.  Faraday's  method  was  to  select  some 
solid  which,  on  being  heated,  gives  off  a  gas.  Some  of  the  solid 
was  placed  in  one  limb,  A,  of  a  bent  tube  (Fig.  32),  the  tube 
was  then  sealed,  and  the  other  limb,  .5,  was  placed  in  ice  water. 
When  the  solid  was  heated,  the  gas  was  given  off  in  a  closed 
space,  and  the  pressure  which  it  exerted  liquefied  a  portion  of 
the  gas  in  the  cold  part  of  the  tube. 


FIG.  32.    Faraday's  method  of 
liquefying  gases 


58 


FIRST  COURSE  IN  CHEMISTRY 


Liquid-air  machines.  Later  it  was  found  that  to  liquefy 
any  given  gas  it  must  be  cooled  below  a  certain  tempera- 
ture, which  varies  from  gas  to  gas  and  which  is  known  as 
the  critical  temperature  of  the  gas.  Machines  are  now  made 
which  simultaneously  compress  the  gas  and  cool  it,  and  all 
gases,  even  air  and  hydrogen,  have  now  been  liquefied. 

Amorphous  and  crystalline 
matter.  Sometimes  the  par- 
ticles of  which  a  piece  of 
solid  matter  is  composed  have 
no  definite  form  and,  even 
under  the  microscope,  have 
no  sharp  edges  or  flat  sur- 
faces. Such  solids  are  said 
to  be  amorphous. 

More  often  a  careful  exam- 
ination of  a  solid  will  show 
that  it  is  made  up  of  a  great 
many  particles,  each  of  which 
has  sharp  edges  and  flat  sur- 
faces. Such  solids  are  said 
to  be  crystalline,  and  each 
individual  piece  is  called  a 
crystal.  While  crystals  have 

a  great  variety  of  forms,  yet  for  any  given  substance  the 
crystalline  form  is  perfectly  definite.  Crystals  range  in 
size  from  microscopic  to  very  large,  a  single  quartz  crystal 
found  in  California  weighing  over  a  ton. 

Although  there  are  a  great  many  forms  of  crystals,  they 
can  all  be  considered  as  varieties  of  only  a  few  fundamental 
ones,  and  a  study  of  the  relations  of  all  these  forms  to  the 
fundamental  ones  constitutes  the  science  of  crystallography. 
In  describing  crystals  wre  shall  not  attempt  to  employ  the 


FIG.  33.   Faraday  (1794-1867) 

An  English  scientist,  who  devised 
methods  for  liquefying  gases 


MATTER  AND  ENERGY  59 

terms  of  crystallography,  but  shall  use  such  terms  as  needle- 
shaped,  flat  plates,  cubes,  and  octahedra.  A  few  well-formed 
crystals  are  shown  in  Fig.  34. 

Heat.  Now  that  we  have  considered  two  elements,  oxy- 
gen and  hydrogen,  and  the  compounds  which  they  form 
with  each  other,  the  question  naturally  arises,  What  is  the 
source  of  the  heat  which  is  given  off  when  these  elements 
combine,  what  is  its  nature,  and  how  can  it  be  meas- 
ured ?  Two  facts  which  have  been  demonstrated  by  experi- 
ment are  of  great  importance  in  answering  these  questions. 


FIG.  34.    Some  examples  of  well-formed  crystals 

Heat  has  no  weight.  One  of  the  most  striking  charac- 
teristics of  heat  is  that  it  has  no  weight,  in  which  respect 
it  differs  from  matter  of  any  kind.  After  a  chemical  action 
in  which  a  great  deal  of  heat  is  set  free,  the  products  weigh 
just  the  same  as  the  original  materials.  This  can  be  shown 
by  causing  the  reaction  to  take  place  in  a  sealed  vessel  and 
weighing  it  before  and  after  the  action. 

Transformations  of  heat.  A  second  remarkable  fact  about 
heat  is  that  it  can  be  changed  into  other  things,  such  as  the 
motion  of  bodies  and  the  electric  current,  and  that  these  in 
turn  may  be  changed  again  into  heat.  Thus,  the  heat  of  the 
flame  A  (Fig.  35)  may  be  transformed  into  motion  by  the 
engine  B.  The  motion  of  the  engine  may  be  used  to  gen- 
erate an  electric  current  in  the  dynamo  C,  and  this  current 


60 


FIRST  COURSE  IN  CHEMISTRY 


may  produce  light  and  heat  in  the  lamp  D  and  also  bring 
about  chemical  decomposition  of  water  into  oxygen  and 
hydrogen  in  the  cell  E. 

Heat  a  form  of  energy.  All  of  these  things  —  heat,  me- 
chanical motion,  electric  current,  and  light  —  are  called 
energy.  The  mechanical  energy  of  moving  bodies  is  con- 
stantly employed  in  doing  work,  and  anything  that  can  do 
work  is  said  to  possess  energy.  Since  all  other  forms  of 


FIG.  35.    Diagram  illustrating  the  transformations  of  heat 

energy  can  be  transformed  into  mechanical  energy,  we  may 
say  that  energy  is  the  capacity  for  doing  work.  Heat  is  there- 
fore one  of  theforrnlTof~energy. 

Conservation  of  energy.  Most  careful  experiment  has 
shown  that  while  energy  can  be  changed  from  one  form  into 
another,  or  transferred  from  one  body  to  another,  the  total 
quantity  remains  the  same.  This  statement  is  known  as  the 
law  of  conservation  of  energy.  For  example,  a  definite 
quantity  of  electrical  energy  will  give  a  definite  quantity 
of  heat  or  power,  otherwise  we  could  not  agree  upon  a 
price  to  pay  for  it.  Conversely,  when  we  find  heat  appear- 
ing under  any  circumstances,  we  may  be  quite  sure  that 


MATTER  AND  ENERGY  61 

energy  in  some  other  form  has  been  used  up.  Thus,  a  hot 
box  on  a  car  is  at  the  expense  of  work  done  by  the  engine. 

Chemical  energy.  When  oxygen  and  hydrogen  combine, 
a  great  deal  of  heat  is  given  off,  yet  there  is  no  very  evi- 
dent form  of  energy  which  is  used  up.  Likewise,  when  the 
electric  current  decomposes  water  into  oxygen  and  hydrogen, 
electrical  energy  is  used  up,  and  our  experience  tells  us  that 
it  must  be  converted  into  some  other  form  of  energy.  We 
assume  that  oxygen  and  hydrogen  possess  some  kind  of 
energy  which  is  changed  into  heat  when  they  unite.  When 
they  are  parted  again,  this  energy  is  restored  at  the  expense 
of  electrical  energy  or  heat.  This  form  of  energy  is  called 
chemical  energy.  It  is  this  form  of  energy  which  all  fuels 
possess  and  for  which  we  pay  when  we  purchase  coal  or 
wood  or  gas. 

Chemical  action.  When  the  chemical  energy  of  a  sub- 
stance is  changed,  we  say  that  chemical  action  has  taken 
place.  It  is  not  always  easy  to  be  sure  that  this  has  oc- 
curred, for  the  change  in  temperature  which  we  observe  may 
be  due  to  the  conversion  of  some  other  kind  of  energy  into 
heat.  If,  however,  a  new  substance  is  formed,  we 'may  be 
sure  that  chemical  action  has  taken  place  in  its  formation. 
(Review  page  9.) 

Unit  of  heat.  Before  a  method  for  measuring  heat  can 
be  devised,  it  is  necessary  to  fix  upon  a  unit  by  which  it 
can  be  measured.  The  unit  of  heat  is  called  a  calorie.  The 
calorie  is  defined  as  the  quantity  of  heat  required  to  raise  the 
temperature  of  one  gram  of  water  through  one  degree. 

Measurement  of  heat.  Matter  in  any  form  can  be  weighed 
on  a  balance,  but  the  measurement  of  heat  is  not  so  simple. 
It  is  accomplished  by  causing  the  heat  to  be  used  up  in  rais- 
ing the  temperature  of  a  weighed  quantity  of  pure  water. 
The  apparatus  in  which  this  is  done  is  called  a  calorimeter 


62 


FIRST  COURSE  IN  CHEMISTRY 


FIG.  36.    A  calorimeter 


and  is  represented  in  Fig.  36.    The  reaction  takes  place  in 

solution  in  the  vessel  A,  and  the  heat  given  off  is  used  up 
in  raising  the  temperature  of  the 
water  contained  in  it,  the  rise  being 
read  on  the  thermometers  D,  D. 
The  water  in  B  serves  to  prevent 
absorption  of  heat  from  the  air. 

Ozone.  Sometimes  chemical  energy 
may  be  added  to  a  substance  in  such 
a  way  as  to  make  a  great  change 
in  its  properties  and  chemical  activ- 
ity while  not  altering  its  percent- 
age composition.  Thus,  if  electric 
sparks  are  passed  through  oxygen 
or  air,  a  small  percentage  of  the 

oxygen  is  converted  into   a   substance    called  ozone.    The 

same  change  can  also  be  brought  about  by  certain  chemical 

processes.     For     example,    if     some 

pieces  of  phosphorus  are  placed  in  a 

bottle    and    partially    covered    with 

water,   the    presence    of    ozone    may 

soon  be  detected  in  the  air  contained 

in  the  bottle. 

The  formation  of  ozone.  The  formation 
of  ozone  may  be  shown  by  partially  cov- 
ering with  water  a  few  pieces  of  stick 
phosphorus  placed  in  the  bottom  of  a 
jar  (Fig.  37).  The  slow  oxidation  of  the 
cold  phosphorus  is  attended  by  the  con- 
version of  some  oxygen  into  ozone.  The 
presence  of  ozone  in  the  air  in  the  jar  is  soon  indicated  by 
its  characteristic  odor,  as  well  as  by  the  property  it  possesses 
of  imparting  a  blue  color  to  strips  of  paper,  A,  previously  dipped 
into  a  solution  of  potassium  iodide  and  starch. 


FIG.  37.   The  formation 
of  ozone  by  the  slow  oxi- 
dation of  phosphorus 


MATTER  AND  ENERGY  63 

The  conversion  of  oxygen  into  ozone  is  attended  by  a 
change  in  volume,  3  volumes  of  oxygen  forming  2  volumes 
of  ozone.  If  the  resulting  ozone  is  heated  to  about  300°, 
the  reverse  change  takes  place,  the  2  volumes  of  ozone 
being  changed  into  3  volumes  of  oxygen.  It  is  possible 
that  traces  of  ozone  exist  in  the  atmosphere,  although  its 
presence  there  has  not  been  definitely  proved. 

Properties.  Ozone  is  a  gas  which  has  the  characteristic 
odor  noticed  about  electrical  machines  when  in  operation. 
When  subjected  to  great  pressure  and  a  low  temperature 
the  gas  condenses  to  a  bluish  liquid,  boiling  at  —  119°.  Its 
chemical  conduct  is  similar  to  that  of  oxygen  except  that 
it  is  far  more  active.  Air  or  oxygen  containing  a  small  per- 
centage of  ozone  is  now  used  in  place  of  oxygen  in  certain 
manufacturing  processes  and  as  a  disinfectant.  It  is  also 
used  to  some  extent  in  the  purification  of  water,  its  strong 
oxidizing  properties  being  sufficient  to  destroy  the  micro- 
organisms present. 

EXERCISES 

1.  Name  three    substances  which  cannot  be    melted  without 
decomposition. 

2.  Name   three    substances  which    cannot    be   boiled   without 
decomposition. 

3.  When  a  pond  begins  to  freeze  in  winter,  why  does  not  all 
the  water  freeze? 

4.  Why  does  a  block  of  ice  melt  so  slowly  even  in  warm  air? 

5.  What  becomes  of  the  heat  applied  to  a  boiling  liquid? 

6.  Why  is  it  necessary  to  boil  eggs  longer  on  a  mountain  top 
than  at  the  seashore  in  order  to  cook  them  ? 

7.  Name  three  crystalline  substances. 

8.  Give  five  illustrations  of  transformation  of  energy  in  daily 
experience. 

9.  What  becomes  of  the  energy  of  burning  coal  in  a  locomotive  ? 
10.    Suggest  three  ways  of  generating  heat.   What  is  the  source  of 

energy  in  each  case  ? 


64  FIRST  COURSE  IN  CHEMISTRY 

11.  Suppose  ice  and  water  to  be  mixed  together  at  0°.    Under 
what  conditions  will  more  water  freeze  ?  more  ice  melt  ? 

12.  Suggest   a  method    of    raising  the  boiling  point  of  water 
above  100°. 

13.  When  water  freezes  in  a  bottle,  why  does  it  break  the  bottle  ? 

14.  Why  is  oil  used  in  machinery? 

15.  How  many  calories  of  heat  are  given  off  in  the  freezing  of 
500  g.  of  water  at  0°?    Ans.  39,500  cal. 

16.  In  a  certain  experiment  2250  g.  of  water  at  20°  was  contained 
in  a  calorimeter.    After  a  reaction  the  temperature  was  at  24.2°. 
How  much  heat  was  evolved?   Ans.  9450  cal. 

TOPICS  FOR  THEMES 

Michael  Faraday  (Thorpe,  Essays  in  Historical  Chemistry). 
Crystals  (see  encyclopedia). 


CHAPTER  VIII 
COMBINING  WEIGHTS ;    THE  ATOMIC  THEORY 

Introduction.  We  have  already  considered  three  laws 
which  deal  with  the  relations  by  weight  which  hold  good 
during  chemical  action:  (1)  the  law  of  conservation  of 
matter,  (2)  the  law  of  definite  composition,  (3)  the  law  of 
multiple  proportion.  To  these  must  now  be  added  a  fourth 
—  the  law  of  combining  weights. 

Combining  weights.  We  have  seen  that  hydrogen  and 
oxygen  combine  in  two  perfectly  definite  ratios  by  weight, 
namely,  1  :  7.94  and  1  :  2  x  7.94.  In  a  similar  way  it  is 
easy  to  determine  the  ratios  in  which  elements  other  than 
oxygen  combine  with  hydrogen.  For  example,  hydrogen 
combines  with  sulfur  to  form  a  gas  called  hydrogen  sulfide, 
and  with  the  metal  calcium  to  form  a  solid  called  calcium 
hydride.  In  these  compounds  the  ratios  by  weight  are 

Hydrogen  sulfide hydrogen  1,  sulfur  16 

Calcium  hydride hydrogen  1,  calcium  19.88 

It  is  therefore  possible  to  assign  to  each  element  combining 
with  hydrogen  a  number  which  expresses  the  weight  in  grams 
of  the  element  which  combines  with  1  g.  of  hydrogen. 

Now  experiment  reveals  a  very  interesting  fact.  The  num- 
bers which  express  the  ratios  in  which  two  elements  com- 
bine with  a  fixed  weight  of  hydrogen  also  express  the  ratio 
in  which  they  combine  with  each  other.  Thus : 

7.1)1  g.  of  oxygen  combines  with  10.88  g.  of  calcium 
2  x  7.94  g.  of  oxygen  combines  with  16  g.  of  sulfur 
19.88  g.  of  calcium  combines  with  16  g.  of  sulfur. 
65 


66 


FIRST  COURSE  IN  CHEMISTRY 


O 

7.94 


S 
16.00 


H 
WO 


Ca 

19.88 


It  is  possible,  therefore,  to  assign  to  each  element  a  number 
which  will  express  the  relative  weight  by  which  it  enters 
into  combination  with  other  elements.  These  numbers  are 
called  the  combining  weights  of  the  elements.  In  Fig.  38 
the  lines  connecting  any  two  symbols  at  once  indicate  the 
ratio  by  weight  in  which  the  elements  combine. 

Elements  with  more  than 
one  combining  weight.  It  is 
evident  that  some  elements 
have  more  than  one  combining 
weight,  for  we  have  seen  that 
1  part  of  hydrogen  combines 
either  with  7.94  or  with  15.88 
parts  of  oxygen.  In  all  such 
cases  the  number  expressing  the 
larger  combining  weight  is  a 
simyl.e  multiple  of  the  number 
expressing  the  smaller  one. 
Standard  for  combining  weights.  The  combining  weights 
are  all  relative  to  some  one  chosen  standard.  It  is  therefore 
possible  to  select  any  element  as  the  standard,  and  any 
desired  weight  of  that  element  as  a  unit  for  comparison. 
For  many  reasons  it  is  better  to  select  oxygen  rather  than 
hydrogen  as  the  standard  element.  It  is  likewise  better  to 
select  8  g.  rather  than  1  g.  as  its  standard  value,  so  that  no 
other  element  may  have  a  combining  weight  of  less  than 
unity.  If  oxygen  is  taken  as  8,  hydrogen  becomes  1.008 ; 
calcium,  20.07;  and  sulfur,  16.03. 

The  law  of  combining  weights.  The  law  of  combining 
weights  may  now  be  stated  as  follows:  To  each  element 
may  be  assigned  a  number  which  in  itself,  or  when  multiplied 
by  some  integer,  expresses  the  weight  by  which  the  element 
combines  with  other  elements. 


FIG.  38.    Diagram   showing    the 
combining  ratios  of  oxygen,  hy- 
drogen, sulfur,  and  calcium 


COMBINING  WEIGHTS;  THE  ATOMIC   THEORY   67 


Natural  laws.  In  science  a  law  is  simply  a  statement  of 
what  might  be  called  a  habit  of  nature.  The  four  laws 
mentioned  at  the  beginning  of  the  chapter  are  merely 
brief  descriptions  of  how  nature  has  been  found  to  act  in 
the  matter  of  chemical  combination.  They  do  not  explain 
anything ;  neither  do  they  compel  chemical  action  to  take 
place  iii  this  way. 

Theories.  Having  formu- 
lated a  number  of  laws 
such  as  these,  we  can  hardly 
help  wondering  why  nature 
should  act  in  so  curious  and 
interesting  a  fashion,  and  we 
set  ourselves  to  imagine  how 
matter  might  be  made  up  so 
that  these  very  laws  would 
be  a  necessary  result.  We 
call  this  process  of  mind 
forming  a  theory. 

The  atomic  theory.  Of 
all  the  theories  that  have 
been  advanced  concerning 
the  nature  of  matter,  the 

one  proposed  by  John  Dalton  (Fig.  39),  and  known  as 
the  atomic  theory,  is  the  most  satisfactory.  The  main  points 
of  this  theory  in  its  present  form,  together  with  the  reasons 
for  making  them  a  part  of  the  theory,  are  as  follows : 

1.  Every    weighable    quantity    of    an    elementary    sub- 
stance is  made  up  of  a  very  great  number  of  unit  bodies 
called   atoms. 

2.  Experiment  shows  that  the  composition  of  a  given 
compound  is  always  the  same.    The  simplest  way  to  adjust 
the  theory  to  this  fact  is  to  assume  that  the  atoms  of  each 


FIG.  39.   John  Dalton  (English) 
(1766-1844) 

The  founder  of  the  atomic  theory 


68  FIRST  COURSE  IN  CHEMISTRY 

element  all  have  the  same  weight,  while  those  of  different 
elements  have  different  weights,  and  that  during  chemical 
union  a  definite  number  of  one  kind  of  atoms  combines 
with  a  definite  number  of  another  kind  to  form  a  particle 
of  a  compound.  If  this  should  be  true,  a  given  compound 
would  of  necessity  have  a  perfectly  definite  composition. 

3.  Since  there  is  no  change  in  weight  when  two  sub- 
stances  act  upon   each   other,   it   must  be   true   that   the 
weights  of  the  individual  atoms  are  unchanged  as  a  result 
of  the  action. 

4.  To  account  for  the  law  of  multiple  proportion  we 
must  assume  that  the  atoms  of  two  different  elements  may 
unite  in  different  ratios ;  for  example,  if  one  atom  of  A 
unites  with  one  of  B  under  one  set  of  conditions,  but  with 
two  of  B  under  other  conditions,  then  we  shall  have  two 
different  compounds.     The  masses  of  B  combined  with  a 
fixed  mass  of  A  will  be  in  the  ratio  of  1  :  2,  since  the 
number  of  atoms  are  in  this  ratio. 

5.  The  law  of  combining  weights  tells  us  that  a  definite 
number  can  be  assigned  to  each  element,  which  expresses 
its  combining  value.     If  each  atom  has  its  own  peculiar 
weight,  and  if  atoms  always  combine  with  each  other  in 
definite  numbers,  then  these  combining  numbers  indicate 
the  relative  weights  of  the  atoms  themselves.    That  an  element 
may  have  two  different  combining  weights,  one  a  multiple 
of  the  other,  is  provided  for  by  the  supposition  that  the 
atoms  are  able  to  combine  in  several  different  ratios. 

Summary  of  the  atomic  theory.  The  atomic  theory  suggests 
that  all  matter  is  made  up  of  minute  bodies  called  atoms. 
The  atoms  of  a  given  element  are  all  alike  in  weight,  but 
those  of  different  elements  have  different  weights.  When 
elements  act  upon  each  other,  the  action  takes  place  between 
definite  small  numbers  of  the  atoms. 


COMBINING  WEIGHTS;  THE  ATOMIC  THEORY  69 

Molecules  and  atoms.  Dalton  applied  the  name  atom  to 
both  elements  and  compounds.  It  is  evident,  however,  that 
the  smallest  particle  of  a  compound  must  consist  of  at  least 
two  different  kinds  of  atoms.  The  term  molecule  is  now 
applied  to  the  smallest  particle  which,  taken  in  large  num- 
bers, makes  up  the  bodies  we  deal  with ;  for  example,  gases 
are  made  up  of  molecules  which  are  moving  about  with 
great  velocity.  The  term  atom  is  applied  to  the  smallest 
unit  of  an  element  which  takes  part  in  a  chemical  action. 

Molecules  of  elements.  Since  two  kinds  of  atoms  unite 
to  form  a  molecule  of  a  compound,  the  question  naturally 
arises,  May  not  two  or  more  atoms  of  the  same  kind  com- 
bine to  form  a  molecule  of  an  elementary  substance  ?  It 
has  been  found  that  the  elements  differ  among  themselves 
in  this  respect.  In  some  cases  the  atoms  do  not  unite ; 
in  other  cases,  as  with  oxygen  and  hydrogen,  two  atoms 
unite  to  form  a  molecule  of  the  element.  The  molecule  of 
ozone  contains  three  atoms.  The  experiments  which  prove 
that  this  is  true  will  be  described  later  (p.  189). 

Value  of  a  theory.  The  value  of  a  theory  is  twofold :  it 
makes  the  processes  of  nature  more  vivid  to  us,  because  it  pre- 
sents them  as  a  picture  rather  than  as  abstract  laws ;  it  also 
leads  to  new  discoveries.  In  adapting  itself  to  known  facts 
and  laws  almost  any  good  theory  will  suggest  a  number  of 
consequences  which  have  not  been  observed,  and  experiments 
can  then  be  made  to  see  whether  these  are  really  as  the  theory 
predicts.  Thus,  the  atomic  theory  predicted  many  properties 
of  gases  which  have  since  been  verified. 

Sometimes  the  progress  of  discovery  will  show  that  a  theory 
unquestionably  expresses  the  truth,  and  the  theory  then  be- 
comes a  statement  of  facts.  There  is  so  much  evidence,  drawn 
from  so  many  sources,  to  show  that  the  atomic  theory  ex- 
presses the  facts  about  the  constitution  of  matter,  that  there 
is  little  doubt  as  to  the  reality  of  atoms  and  molecules. 


TO  FIRST  COURSE  IN  CHEMISTRY 

Atomic  weights.  It  would  be  of  great  interest  if  we  could 
determine  the  weights  of  the  various  kinds  of  atoms.  They 
are  so  very  small,  however,  that  we  can  never  hope  to  deter- 
mine their  weight  even  approximately.  It  has  been  shown 
that  the  smallest  particle  visible  with  the  most  powerful 
microscope  ever  constructed  contains  at  least  1000  atoms. 

We  have  seen,  however,  that  the  ratio  between  the 
combining  weights  is  the  same  as  between  the  weights  of 
the  atoms  themselves,  so  that  we  should  be  able  to  deter- 
mine their  relative  weights  with  precision.  But  most  of  the 
elements  have  more  than  one  combining  weight,  and  we 
must  find  some  means  of  choosing  the  one  which  correctly 
expresses  the  weight  of  a  single  atom. 

It  has  been  found  that,  before  this  problem  can  be  solved, 
methods  must  be  devised  for  finding  the  relative  weights 
of  molecules  of  compounds.  Such  methods  have  been  devel- 
oped and  will  be  described  later  on  (Chap.  XX).  These 
methods  have  led  to  the  adoption  of  a  single  number  for 
each  element,  called  its  atomic  weight.  A  list  of  atomic 
weights  will  be  found  on  the  back  cover  of  the  book.  In 
every  case  they  are  either  the  smallest  combining  weight 
or  some  multiple  of  it. 

EXERCISES 

1.  What  other  scientific  laws  can  you  think  of  besides  the  ones 
mentioned  at  the  beginning  of  this  chapter  ? 

2.  Suppose  the  combining  weight  of  oxygen  had  been  chosen 
as  100  instead  of  8  ;  what  would  then  be  the  combining  weight  of 
hydrogen?   Am.  12.6. 

TOPICS  FOR  THEMES 

John  Dalton  (see  encyclopedia). 

The  atoms  of  the  Greek  philosophers  and  the  atom  of  Dalton 
(see  encyclopedia). 


CHAPTER  IX 
FORMULAS;  EQUATIONS;  CALCULATIONS 

Percentage  composition.  Just  as  we  can  determine  the 
composition  of  water  with  great  accuracy  (p.  46),  so,  by 
similar  means,  we  can  determine  the  composition  of  other 
compounds.  Having  analyzed  a  given  compound,  we  usu- 
ally express  its  composition  in  percentages,  or  in  the  parts 
of  each  element  in  100  parts  of  the  compound.  Thus, 
we  have  seen  that  water  consists  of  88.81  per  cent  of 
oxygen  and  11.19  per  cent  of  hydrogen.  This  mode  of 
expression  takes  no  account  of  the  fact  that  compounds 
are  made  up  of  molecules,  the  atoms  of  which  each  have 
characteristic  weights.  It  would  be  much  better  to  have 
a  method  of  stating  composition  which  would  express  all 
these  facts. 

Atomic  composition.  Remembering  that  the  atomic  weight 
of  oxygen  is  16,  it  is  evident  that  if  we  divide  the  percent- 
age of  oxygen  in  water  by  16,  the  quotient  (5.55)  will  be 
the  relative  number  of  oxygen  atoms  in  100  parts  of  water. 
In  like  manner,  if  we  divide  the  percentage  of  hydrogen 
(11.19)  by  the  atomic  weight  of  the  element  (1.008),  the 
quotient  (11.10)  will  express  the  relative  number  of  hydro- 
gen atoms  in  100  parts  of  water.  The  two  numbers,  5.55 
and  11.10,  therefore  represent  the  ratio  between  the  number 
of  oxygen  and  hydrogen  atoms  in  100  g.  of  water.  But  this 
same  ratio  must  hold  for  any  other  quantity  of  water,  even 
for  one  molecule,  since  any  quantity  of  water  is  made  up  of 

71 


72  FIRST  COURSE  IN  CHEMISTRY 

molecules.  To  reduce  the  ratio  to  its  simplest  terms  we 
divide  the  two  numbers  by  the  smaller  one : 

5.55-*- 5.55  =  1;  11.10-*- 5.55  =  2 

The  ratio  of  oxygen  atoms  to  hydrogen  atoms  in  a  molecule 
of  water  is  therefore  1 :  2. 

Formulas.  We  may  express  the  ratio  found  for  the  oxy- 
gen and  hydrogen  atoms  in  a  molecule  of  water  by  writing 
the  two  symbols  together  thus,  H2O,  the  subscript  2  indi- 
cating that  two  atoms  of  hydrogen  are  in  combination  with 
one  of  oxygen.  This  is  known  as  the  formula  of  water. 

Formula  of  potassium  chlorate.  In  like  manner  we  have  found 
that  potassium  chlorate  consists  of  31.9  per  cent  potassium, 

28.9  per  cent  chlorine,  and  39.2  per  cent  oxygen. 

i 
31.9    -4-  39.2  =  0.8161  =  atomic  weights  of  potassium  in  100  g. 

potassium  chlorate^ 
28.92  -j-  35.45  =  0.8159  =  atomic  weights  of  chlorine  in  100  g. 

potassium  chlorate 

39.16  -f-  16  =  2.447    =  atomic  weights   of    oxygen  in  100  g. 
potassium  chlorate 

Dividing  the  three  quotients  by  the  smallest  (0.8159),  we  get  the 
integers  1, 1,  3.  The  formula  of  potassium  chlorate  is  therefore 
KC103.  Since  analyses  are  always  slightly  inaccurate,  the  atomic 
ratios  will  usually  differ  slightly  from  integers,  but  there  will 
be  no  doubt  as  to  what  the  integer  should  be  in  a  given  case. 

Facts  expressed  by  formulas.  Formulas  are  used  to  ex- 
press several  distinct  facts : 

1.  Atomic  composition  of  molecules.  A  formula  shows  the 
number  and  kinds  of  atoms  in  a  molecule  of  a  compound. 
The  formula  H2O  states  that  a  molecule  of  water  is  composed 
of  two  atoms  of  hydrogen  and  one  of  oxygen.  The  formula 
of  sulfuric  acid,  H2SO4,  shows  that  its  molecule  consists  of 
two  atoms  of  hydrogen,  one  of  sulfur,  and  four  of  oxygen. 


FORMULAS;  EQUATIONS;  CALCULATIONS      73 

2.  Molecular  weights  of  compounds.    Since  each  atom  has 
its  own  weight,  the  sum  of  all  the  atoms  in  a  molecule  must 
be  the  weight  of  the  molecule  itself  relative  to  oxygen  taken 
as  16.     The  relative  weight  of  the  molecule  of  water  is 
therefore  (2  x  1.008)  +  16  =  18.016.    The  relative  weight 
of  the  molecule  of  sulfuric  acid  is  (2  x  1.008)  +  32.06  + 
(4  x  16)  =  98.076. 

3.  Percentage  composition  of  compounds.   From  the  formula 
of  a  compound  we  can  easily  go  back  to  the  percentages  from 
which  it  was  calculated.  Thus,  if  the  molecule  of  water  weighs 
18.016  and  contains  one  oxygen  atom  of  weight  16,  the  frac- 

"1  £» 

tion  of  its  weight  due  to  oxygen  is  —       — ,  or  88.81  per  cent. 

18.016 

The  fraction  due  to  hydrogen  is  ^        —  =  11.19  per  cent. 

lo.Olb 

The  molecule  of  sulfuric  acid  weighs  98.076.    Of  this, 

2.016  32  06 

(or  2.05  per  cent)  is  hydrogen,  -^  ^- — 7  (or  32.70  per 


98.076  v  '  98.076 

cent)  is  sulfur,  and  -  — — —  (or- 65.25  per  cent)  is  oxygen. 

9o.07b 

Gram-molecular  weights ;  formula  weights.  For  practi- 
cal purposes  we  deal  with  pounds  or  with  grams  of  a 
substance,  not  with  atoms  and  molecules.  Now,  since  the 
numbers  18.016,  16,  and  2.016  represent  the  ratio  by 
weight  between  a  molecule  of  water  and  the  oxygen  and 
hydrogen  of  which  it  is  composed,  the  same  ratios  must 
hold  between  any  weight  of  water  we  may  choose  and  the 
oxygen  and  hydrogen  in  this  weight  of  water.  Evidently, 
in  18.916  Ib.  of  water  there  will  be  16  Ib.  of  oxygen  and 
2.016  Ib.  of  hydrogen,  and  in  18.016  g.  there  will  be  16  g. 
of  oxygen  and  2.016  g.  of  hydrogen. 

For  practical  purposes,  therefore,  we  may  allow  the  sym- 
bol H  to  stand  for  1.008  g.  of  hydrogen,  the  symbol  O 


74  FIRST  COURSE  IN  CHEMISTRY 

for  16  g.  of  oxygen,  and  the  formula  H2O  for  18.016  g.  of 
water.  The  weight  in  grams  of  an  element,  corresponding 
to  its  atomic  weight,  is  called  a  gram-atomic  or  symbol 
weight.  The  weight  in  grams  of  a  compound,  corresponding 
to  its  molecular  weight,  is  called  a  gram-molecular  iveight  or 
formula  weight. 

Equations.  Having  devised  a  convenient  way  of  express- 
ing the  composition  of  compounds,  not  in  percentages  but 
in  formulas,  we  make  use  of  equations  to  express  chemical 
transformations,  using  an  arrow  in  place  of  an  equality 
sign.  For  example,  the  equation 

2  H  +  O *•  H20  (1) 

is  a  concise  method  of  stating  two  distinct  facts. 

1.  Qualitatively,  it  states  that  water  is  formed  by  the 
union  of  hydrogen  and  oxygen. 

2.  Quantitatively,  it  tells  us  that  2  symbol  weights  of  hy- 
drogen (2.016  g.)  combine  with  1  symbol  weight  of  oxygen 
(16  g.)  to  form  a  formula  weight  of  water  (18.016  g.). 

Molecular  equations.  Since  a  formula  expresses  the  com- 
position of  a  molecule,  and  since  experiment  has  shown  that 
a  molecule  of  oxygen  and  one  of  hydrogen  each  contain 
two  atoms,  the  formulas  of  these  gases  are  written  O2  and 
H2  rather  than  2  O  or  2  H,  which  would  simply  represent 
two  atoms  not  combined.  If  we  wish  our  equation  to  state 
these  additional  facts,  we  shall  have  to  change  it  to  the  form 

2H2  +  02— )-2H20  (2) 

This  is  called  a  molecular  equation,  and  it  will  be  seen  that 
it  expresses  the  same  ratios  by  weight  as  does  equation  (1). 
It  also  expresses  the  fact  that  2  molecules  of  hydrogen 
combine  with  1  molecule  of  oxygen  to  form  2  molecules 
of  water,  and  this  makes  it  a  more  useful  equation. 


FORMULAS;  EQUATIONS;  CALCULATIONS      75 

Decomposition  of  potassium  chlorate.  Let  us  take  another  ex- 
ample. It  will  be  remembered  that  oxygen  was  prepared  by 
heating  potassium  chlorate,  which  has  the  formula  KC10g. 
When  heated,  this  compound  decomposes  into  oxygen  and  a 
compound  called  potassium  chloride,  whose  formula  is  KC1. 
The  decomposition  is  represented  by  the  equation 

2KC103 ^2KC1  +  302 

This  equation  states  the  following  facts : 

1.  Qualitatively,  potassium  chlorate  decomposes  into  potas- 
sium chloride  and  oxygen. 

2.  Quantitatively,  2  formula  weights  of  potassium  chlorate 
(2  x  122.6  g.)   decompose  into  2   formula  weights   of  potas- 
sium chloride  (2  x  74.69  g.)  and  3  formula  weights  of  oxygen 
(3  x  32  g.).    The  coefficient  before  a  formula  applies  to  the 
formula  as  a  whole,  while  the  subscript  number  applies  only 
to  the  symbol  which  it  follows. 

3.  Molecular ly,  2  molecules  of  potassium  chlorate  decompose 
into  2  molecules  of  potassium  chloride  and  3  of  oxygen. 

Equations  of  reactions  so  far  studied.  Let  us  now  put 
into  the  form  of  equations  a  number  of  the  reactions  studied 
up  to  this  point,  remembering  that  all  of  these  equations 
rest  upon  careful  experimental  analysis. 

1.  Burning  of  mercury : 

2Hg  +  0,— *2HgO. 

2.  Preparation  of  oxygen  : 
From  mercuric  oxide : 

2  HgO >•  2  Hg  +  O2 

From  potassium  chlorate : 

2  KC1O8  — >-  2  KC1  +  3  O2 
From  the  electrolysis  of  water: 

2H20— »2H2+02 


76  FIRST  COURSE  IN  CHEMISTRY 

3.  Preparation  of  hydrogen : 
From  sodium  and  water: 

2  Na  +  2  H2O >-  2  NaOH  +  H2 

From  zinc  and  sulfuric  acid : 

Zn  +  H2S04 *•  ZnS04  +  Ha 

From  steam  and  iron : 

3Fe  +  4H20 ^Fe8O4  +  4Ha 

4.  Preparation  of  hydrogen  peroxide : 

BaO2  -h  H2SO4 *  BaSO4  +  H2O2 

5.  Preparation  of  ozone : 

80,— ^20, 

Representation  of  the  heat  of  reaction.  We  can  also 
employ  chemical  equations  to  express  the  heat  given  off 
or  absorbed  during  chemical  action.  The  equation 

2  H2  +  O2 >•  2  HaO  +  136,800  cal. 

states  the  fact  that  when  4.032  g.  of  hydrogen  combines 
with  32  g.  of  oxygen,  forming  36.032  g.  of  water,  heat 
is  given  off  to  the  extent  of  138,000  cal.  Evidently, 
when  1  formula  weight  (18.016  g.)  of  water  is  formed, 
69,000  cal.  is  given  off,  and  this  is  called  the  heat  of 
formation  of  water. 

Conditions  of  a  reaction  not  indicated  by  equations.  Equa- 
tions merely  state  the  composition  of  the  substances  taking 
part  in  the  reaction  and  the  weights  of  each  one  involved, 
together  with  the  energy  change  measured  as  heat.  They 
do  not  tell  the  conditions  under  which  the  reaction  will 
take  place.  For  example,  the  equation 

2  HgO *  2  Hg  +  O2 


FORMULAS;  EQUATIONS;  CALCULATIONS      77 

does  not  tell  us  that  it  is  necessary  to  keep  heating  the 
mercuric  oxide  to  a  moderately  high  temperature  in  order 
to  effect  its  decomposition.  The  equation 

Zn  4-  H2S04 >•  ZnS04  +  H2 

in  no  way  indicates  that  the  hydrogen  sulfate  must  be  dis- 
solved in  water  before  it  will  act  upon  zinc.  The  equation 

S  +  02^S02 

does  not  indicate  that  no  perceptible  action  takes  place  un- 
less the  sulfur  is  first  heated,  but  that  when  once  started 
it  goes  on  of  its  own  accord  and  with  a  bright  name. 

It  will  therefore  be  necessary  to  pay  close  attention  to 
the  details  of  the  conditions  under  which  a  given  reaction 
occurs,  as  well  as  to  the  statement  of  the  equation  itself. 

Problems  based  on  equations.  Since  an  equation  is  a 
statement  of  the  weights  of  materials  which  take  part  in  a 
reaction,  when  the  equation  has  once  been  established  by 
experiment  we  can  use  it  in  calculating  the  various  weights. 
A  few  examples  will  show  how  this  may  be  done. 

1.  How  many  grams  of  oxygen  are  evolved  on  heating 
100  g.  of  mercuric  oxide  ? 

First  write  the  equation  for  the  reaction  involved : 

2HgO ^2Hg  +  02  (1) 

Next  determine  the  relative  weights  of  the  amounts  of  the 
different  substances  involved  in  the  reaction.  The  atomic 
weights  of  mercury  and  oxygen  are  respectively  200.6  and  16 
(see  table  on  back  cover).  Hence  the  relative  weight  of  the 
2  HgO  equals  2(200.6  +  16),  or  433.2.  Similarly,  the  relative 
weight  of  the  oxygen  evolved,  namely,  O2,  equals  2  x  16,  or 
32.  It  is  convenient  now  to  write  these  numbers  under  the 
formulas  in  equation  (1).  This  then  becomes 

2HgO ^2Hg  +  02 

433.2  32 


78  FIRST  COURSE  IN  CHEMISTRY 

These  numbers  indicate  that  433.2  units  by  weight  (in  this 
case  grams)  of  mercuric  oxide  will,  on  heating,  evolve  32  units 
by  weight  of  oxygen ;  hence  1  g.  of  mercuric  oxide  will  give 

00  00 

^^  g.  of  oxygen,  and  100  g.  will  give  100  x  ^g^  °r  7-38  g. ; 

or  the  relation  between  the  weights  of  the  substances  involved 
may  be  stated  in  the  form  of  a  proportion : 

433.2  :  32  :  :  100  :  x 
433.2  x  =  3200 
x  =  7.38 

2.  I  wish  to  prepare   100  g.   of    oxygen,   using  potassium 
chlorate  as  a  source  of  the  oxygen.    How  many  grains  of  the 
chlorate  will  be  required  ? 

2  KC103 *•  2  KC1  +  3  02 

245.12  96 

Proportion  :    245.12  :  96  :  :  x  :  100 ;  or  x  =  255.33  g.,  Ans. 

3.  How  many  grains  of  zinc  must  be  dissolved  in  sulfuric 
acid  to  produce  10  g.  of  hydrogen  ? 

Zn  +  H2S04 *  ZnS04  +  Ha 

65.37  2.016 

Proportion  :    65.37  :  2.016  :  :  x  :  10 ;  or  x  =  324.2  g.,  Ans. 

It  must  be  remembered  that  the  equations  show  relations 
by  weight,  not  by  volume ;  hence  in  problems  involving  vol- 
umes of  gases  it  will  be  necessary  to  first  find  the  weights  of 
the  gases.  The  table  in  the  Appendix  gives  the  weight  of  1  liter 
of  each  of  the  common  gases,  measured  under  standard  con- 
ditions. The  following  problem  will  illustrate  the  method : 

4.  How  many  grams  of  potassium  chlorate  are  necessary 
to  prepare  100  liters  of  oxygen  ? 

Since  1  liter  of  oxygen  weighs  1.429  g.,  100  liters  will 
weigh  142.9  g. 

2  KC108 *•  2  KC1  +  3  O2 

245.12  96 

Proportion  :    245.12  :  96  :  :  x  :  142.9  ;  or  x  =  363.8  g.,  Ans. 


FORMULAS;  EQUATIONS;  CALCULATIONS      79 

Suggestions.  In  working  such  problems,  do  not  carry 
divisions  beyond  the  second  decimal  place.  If  the  third 
figure  would  be  above  5,  add  1  to  the  second  decimal 
figure.  Having  completed  such  a  problem,  look  to  see  if 
the  result  is  reasonable. 

EXERCISES 

1.  State  all  the  facts  expressed  by  the  formulas  HC1;  HNO3 ; 
Ca(OH)2;  H3PO4. 

2.  State  all  that  is  implied  in  the  equation 

3  O2 >•  2  O3  -  64,800  cal. 

3.  From  the  following  analyses  calculate  the  simplest  formula: 

(1)  S  =  39.07%      O  =  58.49%     H  =  2.44%       Ans.  H2SO3 

(2)  Ca  =  29.40%      S  =  23.56%      O  =  47.04%     Ans.  CaSO4 

(3)  K  =  38.67%     N  =  13.88%      O  =  47.45%     Ans.  KNO3 

4.  It  is  required  to  prepare  30  g.  of  oxygen  by  heating  mercuric 
oxide.    How  much  oxide  must  be  heated?    Ans.  406.13  g. 

5.  What  weight  of  oxygen  can  be  obtained  by  heating  100  g.  of 
potassium  chlorate?   Ans.  39.17  g.    What  volume  will  this  occupy 
under  standard  conditions?    A ns.  27.41  liters. 

6.  What  weight  of  hydrogen  will  be  obtained  by  acting  upon 
100  g.  of  zinc  with  sulfuric  acid?    Ans.  3.08  g.    What  will  be  its 
volume  under  standard  conditions?    A  ns.  34.3  liters. 

7.  A  given  volume  of  oxygen  standing  over  water  at  20°  and 
745  mm.  measures  10  liters.  What  would  be  its  volume  under  standard 
conditions?    Ans.  8.92  liters.    What  is  its  weight ?    Ans.  12.75  g. 

8.  In  making  15  liters  (standard)  of  hydrogen  from  zinc  and  sul- 
furic acid,  what  weight  of  zinc  sulfate  would  be  formed  ?  Ans.  19.8  g. 


CHAPTER  X 


NITROGEN  AND  THE  RARE  ELEMENTS  :  ARGON,  HELIUM, 
NEON,  KRYPTON,  XENON 

Historical.     Nitrogen    was    discovered  by   the    Scottish 

chemist  Rutherford  in  1772.    A  little  later  the  Swedish 

chemist  Scheele  (Fig.  40) 
showed  it  to  be  a  constituent 
of  air,  and  Lavoisier  gave  it 
the  name  azote,  which  means 
that  it  will  not  support  life. 
The  name  nitrogen  was  after- 
wards given  it  because  of  its 
presence  in  niter. 

Occurrence.  Dry  air  is  com- 
posed principally  of  oxygen 
and  nitrogen  in  the  free  state, 
about  78  parts  out  of  every 
100  parts  by  volume  being 
nitrogen.  Nitrogen  also  oc- 
curs in  nature  in  the  form 
of  potassium  nitrate  (KNOg) 
(commonly  called  saltpeter 
or  niter),  as  well  as  in  sodium 

nitrate  (NaNO3).    It  is  also  an  essential  constituent  of  all 

living  organisms  (refer  to  table,  p.  10). 

Preparation.  Nitrogen  can  be  readily  obtained  either  from 

air  or  by  decomposition  of  compounds  containing  the  element. 

80 


FIG.  40.    Karl  Wilhelm  Scheele 
(1735-1784) 

A  famous  Swedish  chemist 


NITROGEN  AND  THE  RAKE  ELEMENTS        81 

1.  Preparation  from  air.  Nitrogen  differs  from  oxygen  in 
that  it  does  not  combine  very  readily  with  most  elements 
save  at  very  high  temperatures.  This  suggests  a  convenient 
method  for  preparing  it  from  air.  It  is  only  necessary  to 
act  upon  a  confined  quantity  of  air  with  some  substance 
which  combines  with  the  oxygen  but  which  has  no  effect 
upon  the  nitrogen.  The  substances  ordinarily  used  for  this 
purpose  are  either  phosphorus  or  copper.  These  are  chosen 


FIG.  41.    The  preparation  of  nitrogen  by  the  action  of  copper  upon  air 

not  only  because  they  combine  readily  with  oxygen  but  be- 
cause the  oxides  formed  are  solids,  and  on  this  account  can 
be  separated  from  the  residue  of  nitrogen  without  difficulty. 
The  nitrogen  obtained  in  these  ways  is  never  quite  pure, 
but  contains  about  1  per  cent  of  a  mixture  of  gases  (chiefly 
argon)  present  in  the  air.  These,  however,  do  not  materially 
affect  the  properties  of  the  nitrogen. 

Details  of  preparation  of  nitrogen.  The  method  used  for  pre- 
paring nitrogen  from  the  air  by  burning  out  the  oxygen  with 
phosphorus  has  already  been  described  (p.  3). 

When  copper  is  used,  the  metal  is  placed  in  a  tube  A  (Fig.  41) 
and  heated.  Air  is  then  forced  slowly  through  the  tube  by  pour- 
ing water  into  the  bottle  B.  The  oxygen  of  the  air  combines 


82 


FIRST  COURSE  IK  CHEMISTRY 


with  the  hot  copper,  forming  the  black  solid,  copper  oxide  (CuO), 
which  remains  in  the  tube,  while  the  nitrogen  passes  on  and  is 
collected  over  water  in  the  cylinder  C. 

2.  Preparation  from  compounds  of  nitrogen.  Ammonium 
nitrite  (NH4NO2)  serves  as  a  convenient  source  for  pre- 
paring pure  nitrogen.  When  the  compound  is  heated,  the 
change  represented  in  the  following  equation  takes  place  : 


NH4NO2 


2  H2O  +  N2 


FIG.  42.    Tubercles  on  the 
roots  of  bean  plants 


Properties.  Nitrogen  is  similar  to  oxygen  and  hydrogen 
in  that  it  is  a  colorless,  odorless,  and  tasteless  gas.  One 

liter  of  nitrogen  weighs  1.2507  g. 
It  is  almost  insoluble  in  water. 
It  can  be  obtained  in  the  form 
of  a  colorless  liquid  having  a 
boiling  point  of  -195.7°.  At 
-  210.5°  it  solidifies. 

Chemical  conduct.  Under  ordi- 
nary conditions  nitrogen  is  very 
inactive.  It  does  not  easily  com- 
bine with  oxygen,  as  is  evident 
from  the  fact  that  the  air  contains  both  of  these  gases  ; 
nor  does  it  combine  with  anything  else  very  readily,  as 
is  apparent  from  the  fact  that  there  is  so  much  of  it  in 
the  air  and  so  little  in  the  earth's  crust  in  combination 
with  other  elements. 

At  high  temperatures  the  activity  of  nitrogen  greatly  in- 
creases. When  it  is  mixed  with  oxygen  and  strongly  heated, 
a  small  fraction  of  the  two  gases  combine,  but  the  action  is 
always  very  incomplete.  The  best  results  are  obtained  by 
passing  electric  sparks  through  a  mixture  of  the  two  gases 
(or  air)  or  by  causing  the  mixture  to  flow  through  an  elec- 
tric arc.  Under  these  conditions  nitric  oxide  (NO)  forms. 


NITROGEN  AND  THE  BARE  ELEMENTS        83 


Under  similar  conditions  nitrogen  and  hydrogen  combine  to  a 
limited  extent  to  form  the  hydride,  ammonia  (NH3).  When 
nitrogen  is  heated  with  metals,  the  action  is  much  more 
energetic,  particularly  with  magnesium,  titanium,  or  alumin- 
ium. The  resulting  compounds  are  called  nitrides,  just  as 
compounds  of  an  element  with  oxygen  are  called  oxides. 

The  assimilation  of  ni- 
trogen by  plants.  While 
nitrogen  is  an  essential 
constituent  of  both  plants 
and  animals,  yet  with  the 
exception  of  a  few  plants 
none  of  these  organisms 
have  the  power  of  directly 
assimilating  free  nitrogen. 
For, example,  the  nitro- 
gen in  our  bodies  is  taken 
from  our  foods  and  not 
from  the  air  which  we 
inhale.  It  has  long  been 
known,  however,  that  cer- 
tain plants,  chiefly  clover, 
alfalfa,  beans,  and  similar 
plants  belonging  to  the 
natural  order  Legumino- 
sae,  not  only  thrive  in  poor  soil  but  at  the  same  time  en- 
rich it.  It  is  now  known  that  these  plants  obtain  at  least  a 
portion  of  their  nitrogen  from  the  atmosphere.  This  is 
accomplished  by  groups  of  microorganisms  which  are  gath- 
ered in  little  tubercles  on  the  roots  of  the  plants  (Fig.  42). 
These  organisms  convert  the  nitrogen  of  the  air  into  com- 
pounds of  nitrogen ;  some  of  these  are  assimilated  by  the 
plant  and  some  are  left  in  the  soil  and  thus  enrich  it. 


FIG.  43.    Sir  William  Ramsay  (1852-) 

The  discoverer  of  the  rare  elements  present 
in  the  atmosphere 


84  FIRST  COURSE  IN  CHEMISTRY 

Argon,  helium,  neon,  krypton,  xenon.  These  are  rare  elements 
and  occur  in  the  air  in  very  small  quantities.  They  are  similar 
in  that  they  are  all  colorless,  odorless  gases.  They  differ  from 
all  other  known  elements  in  that  they  are  entirely  inert,  form- 
ing no  compounds  whatever.  Argon,  the  most  abundant  of  the 
group,  was  discovered  in  1894  by  two  British  scientists,  Lord 
Rayleigh  and  Sir  William  Ramsay  (Fig.  43).  In  1868  Lockyer 
showed  that  a  gaseous  element,  to  which  he  gave  the  name 
helium,  was  present  in  the  gases  surrounding  the  sun.  In  1895 
Ramsay  showed  that  this  same  element  was  present  in  the 
gases  evolved  in  heating  certain  minerals,  and  later  that  traces 
of  it  were  present  in  the  atmosphere.  The  three  remaining 
members  of  the  group  were  discovered  by  Ramsay  and  Travers 
in  1898.  They  obtained  them  from  liquid  air,  thus  proving 
their  presence  in  the  atmosphere. 

Helium  is  of  interest  in  that,  of  all  known  gases,  it  is  the 
most  difficult  to  liquefy.  It  was  first  obtained  in  liquid  form 
in  1908.  The  boiling  point  of  the  liquid  is  -  268.7°.  Some  facts 
pertaining  to  these  gases  are  given  in  the  following  table : 


TABLE  OF  RARE  ATMOSPHERIC  ELEMENTS 


HELIUM 

NEON 

ARC  ox 

KRYPTON 

XENON 

Weight  of  1  liter    .     . 
Boiling  point  of  liquid 
form      

0.1782g. 
—  268.7° 

0.9002  g. 

-239° 

1.7809g. 
-186° 

3.708g. 
-151.7° 

5.851  g. 
-109° 

Number  of  volumes  in 
1,000,000  volumes  of 
air  (approximate)    . 

4.00 

12.3 

9400 

0.05 

0.006 

EXERCISES 


1.  Why  not  prepare  nitrogen  by  burning  a  candle  in  confined  air  ? 

2.  In  Fig.  4  why  does  the  withdrawal  of  oxygen  cause  the  water 
to  rise  in  the  bell  jar  ? 

3.  Which  contains  the  greater  percentage  of  nitrogen :  sodium 
nitrate  or  potassium  nitrate  ? 


NITROGEN  AND  THE  RARE  ELEMENTS        85 

4.  100  liters   of  dry  air  contain  how  many  liters  of  nitrogen? 
What  is  the  weight  of  this  volume  of  nitrogen?    Ans.  97.55  g. 

5.  What  weight  of  ammonium  nitrite  will  be  required  for  the 
preparation  of  10  g.  of  nitrogen?    Ans.  22.9  g.    What  will  be  the 
volume  of  the  nitrogen  under  standard  conditions  ?    A  ns.  8  liters. 

6.  A  quantity  of  nitrogen  standing  over  water  at  20°  and  750  mm. 
measured  25  liters.    What  will  be  the  volume  under  standard  con- 
ditions? A  ns.  22.9  liters. 

TOPICS  FOR  THEMES 

Argon  and  the  rare  atmospheric  gases  (McPherson  and  Hender- 
son, A  Course  in  General  Chemistry). 

Fixation  of  nitrogen  by  plants  (Duncan,  Chemistry  of  Commerce  ; 
also  write  to  the  Department  of  Agriculture,  Washington,  D.  C., 
for  bulletins). 


CHAPTER  XI 
THE  ATMOSPHERE 

Historical.  The  terms  atmosphere  and  air  are  often  used 
interchangeably,  although  strictly  speaking  the  former  term 
is  applied  to  the  entire  gaseous  envelope  surrounding  the 
earth,  while  the  latter  is  applied  to  a  limited  portion  of 
this  envelope.  Like  water,  air  was  formerly  regarded  as 
an  element.  Near  the  close  of  the  eighteenth  century, 
however,  through  the  experiments  of  Scheele,  Priestley, 
Cavendish,  and  Lavoisier,  it  was  shown  to  be  a  mixture  of 
at  least  two  gases  —  those  which  we  now  call  oxygen  and 
nitrogen.  By  absorbing  the  oxygen  from  an  inclosed  vol- 
ume of  air  and  measuring  the  contraction  in  volume  due 
to  the  removal  of  oxygen,  Cavendish  was  able  to  determine 
with  considerable  accuracy  the  relative  volumes  of  oxygen 
and  nitrogen  present. 

Composition  of  the  air.  The  normal  constituents  of  air, 
together  with  the  approximate  volumes  of  each  in  samples 
collected  in  the  open  fields,  are  as  follows : 

Oxygen 21  volumes  in  100  volumes  of  dry  air 

Nitrogen     .....  78  volumes  in  100  volumes  of  dry  air 

Water  vapor    ....  variable  within  wide  limits 

Carbon  dioxide    .     .     .  3  to  4  volumes  in  10,000  volumes  of  dry  air 

Argon 0.940  volumes  in  100  volumes  of  dry  air 

Helium,  neon,  krypton, 

xenon traces 

In  addition  there  are  usually  present  small  quantities  of  hy- 
drogen peroxide,  ammonium  nitrate,  microorganisms,  dust 

86 


THE  ATMOSPHERE  87 

particles,  and  traces  of  hydrogen.  Although  not  definitely 
proved,  it  is  probable  that  small  amounts  of  ozone  are  also 
present.  The  air  in  large  cities  and  manufacturing  districts 
is  also  likely  to  contain  certain  gases  evolved  in  manufac- 
turing processes.  Among  these  are  hydrogen  sulfide  (H2S) 
and  sulfur  dioxide  (SO2). 

Water  vapor  in  the  air.  The  quantity  of  water  vapor 
which  may  be  present  in  the  air  varies  with  the  tempera- 
ture. This  is  shown  in  the  following  table,  which  gives  the 
weight  in  grams  of  the  water  vapor  that  1  cu.  m.  of  air  can 
absorb  at  the  temperature  indicated : 

Temperature,  0°  10°  20°  30° 

Weight  of  water,         4.8  g.         9.9  g.         17.1  g.         30  g. 

The  constituents  of  the  air  that  are  essential  to  life.  The 
constituents  that  are  known  to  be  essential  to  life  are  oxy- 
gen, nitrogen,  water  vapor,  and  carbon  dioxide.  The  first 
three  of  these  have  already  been  discussed  in  detail.  The 
remaining  one,  carbon  dioxide,  is  a  gas  having  the  formula 
CO2.  It  is  evolved  in  the  processes  of  both  respiration  and 
combustion  (p.  20),  so  that  large  quantities  of  it  are  con- 
stantly being  added  to  the  atmosphere.  The  properties  of 
the  gas  will  be  described  in  the  chapter  relating  to  the 
compounds  of  carbon ;  it  is  only  necessary  to  note  here 
that  it  is  a  comparatively  heavy  gas  and  will  neither  burn 
nor  support  combustion. 

The  oxygen  in  the  atmosphere  directly  supports  life 
through  the  process  of  respiration.  The  nitrogen  serves  to 
dilute  the  oxygen  and  thus  to  diminish  the  intensity  of  its 
action.  It  is  likewise  assimilated  by  certain  plants  (p.  83). 
The  water  vapor  prevents  excessive  evaporation  of  the 
water  present  in  organisms,  while  the  carbon  dioxide  is  an 
essential  plant  food. 


88 


FIRST  COURSE  IN  CHEMISTRY 


The  quantitative  analysis  of  air.  A  number  of  different 
methods  have  been  devised  for  the  determination  of  the 
percentages  of  the  constituents  of  the  atmosphere.  Among 
these  are  the  following: 

1.  Determination  of  oxygen.  The  oxygen  is  withdrawn  from 
a  measured  volume  of  air  inclosed  in  a  tube,  by  means  of 

phosphorus. 

To  make  the  determination,  a 
graduated  tube  is  filled  with  water 
and  inverted  in  a  vessel  of  water. 
A  sample  of  the  air  to  be  analyzed 
is  then  introduced  into  the  tube 
until  it  is  nearly  filled  with  the  gas, 
and  the  volume  is  carefully  noted. 
A  small  piece  of  phosphorus  is  at- 
tached to  a  wire  and  brought  within 
the  tube  as  shown  in  Fig.  44.  After 
a  few  hours  the  oxygen  in  the  in- 
closed air  will  have  combined  with 
the  phosphorus,  the  water  rising  to 
take  its  place.  The  phosphorus  is 
removed,  and  the  volume  is  again 

noted.   The  contraction  in  the  volume  of  the  air  is  equal  to 

the  volume  of  oxygen  absorbed. 

2.  Determination  of  nitrogen.    If   the  gas  left  after  the 
removal  of  oxygen  from  a  portion  of  air  is  passed  over 
heated    magnesium,    the    nitrogen   is    withdrawn,    leaving 
argon  and  the  other  rare  elements.    It  may  thus  be  shown 
that,  of  the  79  volumes  of  gas  left  after  the  removal  of  the 
oxygen  from   100   volumes   of  air,   approximately   78   are 
nitrogen  and  0.93  argon.    The  other  elements  are  present 
in  such  small  quantities  that  they  may  be  neglected. 

3.  Determination  of  water  vapor  and  carbon  dioxide.    These 
constituents  are  determined  by  passing  a  known  volume  of 


FIG.  44.  The  withdrawal 
of  oxygen  from  a  measured 
volume  of  air  by  phosphorus 


THE  ATMOSPHERE  89 

air  through  two  tubes,  the  first  containing  calcium  chlo- 
ride, and  the  second  calcium  hydroxide,  or,  better,  sodium 
hydroxide.  The  calcium  chloride  removes  the  moisture, 
while  the  sodium  hydroxide  removes  the  carbon  dioxide. 
The  increase  in  the  weights  of  these  two  substances  will 
give  the  weights  of  moisture  and  carbon  dioxide  respectively 
in  the  original  volume  of  air. 

Processes  tending  to  change  the  composition  of  the  air. 
These  processes  naturally  fall  into  two  classes :  those  which 
increase  the  carbon  dioxide  and  those  which  diminish  it. 

1.  Processes  tending  to  increase  the  quantity  of  carbon  dioxide. 
Not  only  do  large  quantities  of  carbon  dioxide  escape  into 
the  atmosphere  from  volcanoes  and  crevices  in  the  earth's 
crust,  but  certain  processes  are  constantly  taking  place  which 
are  attended  by  evolution  of  this  gas.    Chief  among  these 
are  the  following :   (a)  Respiration.    In  this  process  some 
of  the  oxygen  in  the  inhaled  air  is  absorbed  by  the  blood 
and  carried  to  all  parts  of  the  body,  where  it  combines  with 
the  carbon  of  the  worn-out  tissues.    The  products  of  oxida- 
tion are  carried  back  to  the  lungs  and  exhaled  largely  in 
the  form  of  carbon  dioxide,    (b)  Combustion.   All  the  ordi- 
nary fuels  contain  large  percentages  of  carbon.   On  burning, 
this  is  oxidized  to  carbon  dioxide,    (c)  Decay  of  organic 
matter.    When  organic  matter  decays  in  the  air  the  carbon 
present  is  oxidized  to  carbon  dioxide. 

2.  Processes  tending  to  decrease  the  quantity  of  carbon  dioxide. 
There  are  two  general  processes  which  tend  to  diminish  the 
quantity  of  carbon  dioxide  in  the  atmosphere. 

(a)  The  action  of  plants.  Plants  have  the  power,  when 
growing  in  sunlight,  of  absorbing  carbon  dioxide  from  the 
air,  retaining  the  carbon  and  returning  a  portion  of  the 
oxygen  to  the  air.  It  is  from  this  source  that  plants  obtain 
their  entire  supply  of  carbon. 


90 


FIRST  COURSE  IN  CHEMISTRY 


That  plants  evolve  oxygen  in  the  sunlight  may  be  shown 
as  follows :  Some  freshly  gathered  leaves  are  placed  under 
water  in  the  jar  A  (Fig.  45)  and  covered  with  the  funnel  B, 
the  stem  of  which  extends  into  the  graduated  tube  C.  Bubbles 
of  oxygen  make  their  escape  from  the  surface  of  the  leaves, 
and  may  be  collected  in  the  measuring  tube  C. 

(b)  The  weathering  of  rocks.  Large  quantities  of  carbon 
dioxide  are  being  constantly  withdrawn  from  the  atmosphere 

through   its  combination  with 
various  rock  materials. 

The  composition  of  the  air 
constant.  Notwithstanding  the 
changes  constantly  taking  place 
which  tend  to  alter  the  compo- 
sition of  the  air,  the  results  of  a 
great  many  analyses  of  air  col- 
lected in  the  open  fields  show 
that  the  percentages  of  oxygen 
and  nitrogen,  as  well  as  of 
carbon  dioxide,  are  very  nearly 
constant.  Indeed,  so  constant 
are  the  percentages  of  oxygen 
and  nitrogen  that  the  question 
has  arisen,  whether  air  is  not 
a  definite  chemical  compound. 

Air  a  mixture.  That  the  oxygen  and  nitrogen  in  the  air 
are  not  combined  may  be  shown  in  a  number  of  ways,  among 
which  are  the  following : 

1.  When  air  dissolves  in  water  it  has  been  found  that 
the  ratio  of  oxygen  to  nitrogen  in  the  dissolved  air  is  no 
longer  21  :  78,  but  more  nearly  35  :  65.  If  air  were  a  chem- 
ical compound,  the  ratio  of  oxygen  to  nitrogen  would  not 
be  changed  by  solution  in  water. 


FIG.  45.  The  liberation  of  oxygen 
from  plants  exposed  to  sunlight 


THE  ATMOSPHERE  91 

2.  A  chemical  compound  in  the  form  of  a  liquid  has  a 
definite  boiling  point  at  a  given  pressure  (p.  57).  Water, 
for  example,  boils  at  100°  under  standard  pressure.  More- 
over, the  steam  which  is  formed  has  the  same  composition 
as  the  water.  The  boiling  point  of  liquid  air,  on  the  other 
hand,  gradually  rises  as  the  liquid  boils,  the  nitrogen  escap- 
ing first  followed  by  the  oxygen.  If  the  two  were  combined, 
they  would  pass  off  together  in  the  ratio  in  which  they  are 
found  in  the  air. 

Why  the  air  has  a  constant  composition.  If  air  is  a  mixture 
and  changes  are  constantly  taking  place  which  tend  to  modify 
its  composition,  how,  then,  do  we  account  for  the  constancy  of 
composition  which  the  analyses  reveal  ?  This  is  explained  by 
several  facts :  (1)  The  changes  which  are  caused  by  the  proc- 
esses of  combustion,  respiration,  and  decay,  on  the  one  hand, 
and  the  action  of  plants,  on  the  other,  tend  to  equalize  each 
other.  (2)  The  winds  keep  the  air  in  constant  motion  and  so 
prevent  local  changes.  (3)  The  volume  of  air  is  so  vast  and  the 
changes  which  occur  are  so  small,  compared  with  the  total  vol- 
ume, that  they  cannot  be  readily  detected.  (4)  Finally,  it  must 
be  noted  that  only  air  collected  in  the  open  fields  shows  this 
constancy  in  composition.  The  air  in  a  poorly  ventilated  room 
occupied  by  a  number  of  people  rapidly  changes  in  composition. 

Impure  air  and  ventilation.  The  difference  in  the  per- 
centages of  oxygen,  carbon  dioxide,  and  moisture  present 
in  inhaled  and  exhaled  air  are  shown  in  the  following  table  : 


CONSTITUENT 

IXHALKD  Alii 

EXHALED  AIK 

Oxvrrpn 

°1  00% 

16  00% 

Carbon  dioxide     

0  04% 

4  38% 

Water  vapor    

variable 

s&rt/ur&tecl 

The  injurious  effects  resulting  from  inadequate  ventilation 
seem  to  be  due  neither  to  lack  of  oxygen  nor  to  the  excess 


92  FIRST  COURSE  IK  CHEMISTRY 

of  carbon  dioxide  ;  rather  they  are  due  to  high  temperature 
and  to  the  presence  of  an  abnormal  amount  of  water  vapor, 
both  of  which  conditions  are  apt  to  prevail  in  crowded 
and  poorly  ventilated  rooms. 

Not  only  is  water  vapor  exhaled  from  the  lungs,  but  there 
is  constant  evaporation  of  moisture  from  the  pores  of  the 
skin,  and  in  this  process  much  heat  is  absorbed.  Notwith- 
standing the  extreme  changes  in  the  temperature  of  the  air, 
the  temperature  of  the  body  in  health  remains  nearly  con- 
stant. It  is  partly  by  variations  in  the  amount  of  mois- 
ture evaporating  from  the  skin  that  the  temperature  of  the 
body  is  maintained  at  this  constant  value.  If  an  abnormal 
amount  of  water  vapor  is  present  in  the  air,  the  evaporation 
of  moisture  from  the  skin  takes  place  very  slowly,  and  bodily 
discomfort  follows.  Moreover,  when  the  air  is  perfectly  still, 
that  portion  of  the  air  in  contact  with  the  body  tends  to  be- 
come saturated  with  moisture,  and  evaporation  diminishes ; 
hence  the  relief  that  comes  from  keeping  the  air  in  motion, 
as  with  an  electric  fan. 

In  general,  a  moisture  content  of  about  70  per  cent  of 
that  required  for  saturation  is  most  conducive  to  comfort. 
The  volume  of  fresh  air  necessary  for  good  ventilation  varies 
greatly  with  conditions,  but  in  general  may  be  said  to  be 
about  30  cu.  ft.  per  minute  for  each  person  present. 

The  properties  of  air.  Inasmuch  as  air  is  composed 
principally  of  a  mixture  of  oxygen  and  nitrogen,  which  ele- 
ments have  already  been  discussed,  its  properties  may  be 
inferred  largely  from  those  of  the  two  gases.  One  liter 
weighs  1.2928  g. 

Liquid  air.  Like  all  other  gases,  air  can  be  liquefied  and 
solidified  by  the  combined  effect  of  pressure  and  low  temper- 
ature. Air  is  now  liquefied,  on  a  commercial  scale  by  com- 
pressing it  by  means  of  powerful  pumps,  the  temperature 


THE  ATMOSPHERE 


93 


FIG.  46.    A  Dewar 
flask   for   preserv- 
ing liquid  air 


of  the  compressed  gas  being  lowered  sufficiently  by  allow- 
ing a  portion  of  the  gas  to  expand  in  such  a  way  that  the 
heat  absorbed  in  the  expansion  is  largely 
withdrawn  from  the  remaining  gas. 

Liquid  air  is  essentially  a  mixture  of 
liquid  nitrogen  (boiling  point,  —  195.7°) 
and  liquid  oxygen  (boiling point,  —182.9°); 
hence  if  liquid  air  is  allowed  to  evapo- 
rate, the  nitrogen  tends  to  vaporize  first. 
Advantage  is  taken  of  this  difference  in 
boiling  points  to  separate  the  oxygen  and 
nitrogen  from  each  other,  and  the  method 
serves  as  a  commercial  one  for  obtaining 
the  two  gases.  Liquid  air  is  also  employed  when  very  low 
temperatures  are  desired. 

Dewar  flasks ;  thermos  bottles.  Liquid  air  may  be  kept 
for  some  hours  in  a  special  form  of  flask  devised  by  the 
Scottish  scientist,  Dewar,  known  as  a  Dewar 
flask.  This  consists  of  two  concentric  ves- 
sels (Fig.  46)  of  any  convenient  shape. 
These  are  joined  together  at  the  upper  rim 
only,  and  the  space  between  them  is  ex- 
hausted by  an  air  pump.  The  vacuum 
serves  as  the  best  possible  insulator  to  pre- 
vent heat  conduction.  The  surface  of  the 
outer  flask  is  often  silvered,  in  order  to 
reflect  the  external  heat  and  thus  prevent 
its  absorption.  The  so-called  thermos  bottles 
(Fig.  47)  are  constructed  on  the  same  plan 
and  are  very  effective  for  keeping  liquids 
either  hot  or  cold  for  several  hours. 

Combustion  in  air  and  in  oxygen.    Knowing  the  composi- 
tion of  air  and  the  properties  of  oxygen  and  nitrogen,  one 


FIG.  47.    A  ther- 
mos bottle 


94  FIBST  COUKSE  IN  CHEMISTRY 

can  readily  understand  why  substances  bum  more  readily 
in  pure  oxygen  than  in  air.  Since  combustion  consists  in 
the  union  of  matter  with  oxygen,  it  is  evident  that  the 
speed  of  the  action  will  be  influenced  by  the  amount  of 
oxygen  which  is  in  contact  with  the  burning  body.  When 
combustion  takes  place  in  air,  only  about  one  fifth  as  much 
oxygen  can  come  in  contact  with  the  body  as  in  pure 
oxygen,  and  the  speed  of  the  action  is  therefore  correspond- 
ingly less.  Moreover,  the  speed  increases  with  the  temper- 
ature. When  a  substance  burns  in  air,  much  of  the  heat 
generated  is  spent  in  heating  the  nitrogen  present.  The 
temperature  reached  is  therefore  less,  and  the  speed  of  the 
action  is  correspondingly  slower. 

Dust  explosions.  Experiments  show  that  under  proper  con- 
ditions a  combustible  body  floating  in  the  air  in  the  form  of  a 
powder  will  burn  almost  instantaneously  when  ignited,  since 
the  extent  of  surface  exposed  to  the  oxygen  is  very  great.  The 
result  is  an  explosion.  Flour-mills  have  been  known  to  explode 
with  great  violence,  and  many  of  the  most  powerful  mine  ex- 
plosions are  due  to  the  combustion  of  the  coal  dust  floating 
in  the  air  of  the  mine,  which  in  some  way  becomes  ignited. 

EXERCISES 

1.  When  oxygen  and  nitrogen  are  mixed  in  the  proportion  in 
which  they  exist  in  the  atmosphere,  heat  is  neither  evolved  nor  ab- 
sorbed by  the  process.    What  important  point  does  this  suggest  ? 

2.  How  does  the  air  in  manufacturing  districts  differ  in  com- 
position from  that  in  the  open  fields? 

3.  Can  you  suggest  any  reason  why  the  growth  of  clover  in  a  field 
improves  the  soil  ? 

4.  When  ice  is  placed  in  a  vessel  containing  liquid  air,  the  latter 
boils  violently.    Explain. 

5.  Does  an  electric  fan  lower  the  temperature  of  a  room?  of  an 
individual  in  the  room  ? 

6.  What  is  the  meaning  of  the  word  thermos  ? 


THE  ATMOSPHERE  95 

7.  Assuming  that  dry  wood  contains  40  per  cent  carbon,  all  of 
which  originally  came  from  carbon  dioxide  in  the  air,  what  weight 
of  CO2  would  have  to  be  absorbed  by  a  plant  to  make  500  g.  of 
wood ?  Ans.  733.3  g.   What  volume  would  this  occupy  under  standard 
conditions?  Ans.  370.9  liters. 

8.  Taking  the  volumes  of  the  oxygen  and  nitrogen  in  100  volumes 
of  air  as  21  and  78  respectively,  calculate  the  percentages  of  these 
elements  present  in  the  air  by  weight. 

9.  Would  combustion  be  more  intense  in  liquid  than  in  gase- 
ous air  ? 

10.  A  tube  containing  calcium  chloride  was  found  to  weigh 
30.1293  g.  A  volume  of  air  which  weighed  15.2134  g.  was  passed 
through,  after  which  the  weight  of  the  tube  was  found  to  be  30.3405  g. 
Find  the  percentage  of  moisture  present  in  the  air.  Ans.  1.39  per  cent. 

TOPICS  FOR  THEMES 

The  atmospheric  conditions  leading  to  fog,  rain,  snow,  and  hail. 
Liquid  air  (see  encyclopedia). 


CHAPTER    XII 
SOLUTIONS  AND  IONIZATION 

Definitions.  When  a  solid  is  thoroughly  stirred  through 
a  liquid  the  solid  often  passes  completely  from  sight  as  an 
individual  body,  or  dissolves,  forming  a  solution.  The  liquid 
in  which  the  substance  dissolves  is  called  the  solvent,  while 
the  dissolved  substance  is  the  solute. 

It  often  happens  that  the  solid  does  not  really  dissolve  but 
remains  suspended  in  the  liquid,  rendering  it  cloudy.  This 
is  true  of  clay  shaken  with  water.  Upon  standing  quietly 
for  some  time,  however,  matter  in  suspension  gradually 
settles,  while  matter  in  true  solution  does  not.  The  rapidity 
of  settling  depends  upon  the  size  of  the  solid  particles. 

Saturated  solutions.  On  adding  a  solid  to  a  liquid  in 
small  portions  at  a  time,  it  will  be  found  that  a  point  is 
reached  at  which  the  liquid  will  not,  at  that  temperature, 
dissolve  more  of  the  solid;  the  solid  and  the  solution 
remain  apparently  unchanged  in  contact  with  each  other, 
the  rate  at  which  the  solid  dissolves  being  just  balanced  by 
the  rate  at  which  solid  crystallizes  out.  This  condition 
may  be  described  by  saying  that  solid  and  solution  are  in 
equilibrium  with  each  other.  A  solution  is  said  to  be 
saturated  at  a  given  temperature  when  it  remains  un- 
changed in  concentration  in  contact  with  some  of  the  solid. 
The  weight  of  the  solid  which  will  completely  saturate  a 
definite  volume  of  a  liquid  at  a  given  temperature  is  called 
the  solubility  of  the  substance  at  that  temperature. 

96 


SOLUTIONS  AND  IONIZATION 


97 


Supersaturated  solutions.  Most  solids  are  more  soluble  in  hot 
than  in  cold  liquids,  and  a  liquid,  saturated  at  a  high  tempera- 
ture, usually  deposits  the  excess  of  solute  in  the  form  of 
crystals  as  the  temperature  falls,  maintaining  saturation  at  all 
temperatures.  Sometimes  the  crystals  fail  to  form  as  the  solu- 
tion cools,  especially  if  it  is  not  disturbed  in  any  way.  The 
solution  then  contains  more  of  the  solute  than  is  normally 
present  when  the  solution  is  in 
equilibrium  with  the  solid.  Such 
a  solution  is  said  to  be  supersat- 
urated. When  a  crystal  of  the 
solid  is  added  to  the  supersatu- 
rated solution  the  excess  of  solute 
at  once  crystallizes  out,  the  crys- 
tallization starting  from  the  added 
crystals.  This  may  be  shown  in 
a  striking  way  by  suspending  a 
small  crystal  by  a  thread  in  a 
supersaturated  solution  (Fig.  48). 
The  crystal  grows  rapidly  in  size, 
and  fragments,  breaking  off,  start 
crystallization  at  other  points. 

Classes  of  solutions.  All  gases 
mix  freely  with  each  other  in  all 
proportions,  and  such  mixtures 
may  be  regarded  as  the  solution 
of  one  gas  in  another.  Gases, 

liquids,  and  solids  dissolve  in  liquids,  and  one  solid  fre- 
quently dissolves  in  another.  The  most  familiar  of  these 
classes  are  solutions  of  gases  or  solids  in  liquids. 

Conditions  affecting  solubility.  A  number  of  different  con- 
ditions influence  the  solubility  of  a  substance  in  a  liquid. 

1.  Nature  of  the  solute.  Each  substance  has  its  peculiar 
solubility  just  as  it  has  its  own  odor,  taste,  or  crystalline 
form.  All  substances  may  be  regarded  as  being  soluble  to 


FIG.  48.    The  rapid  growth  of 
a  crystal  suspended  in  a  super- 
saturated solution 


98  FIRST  COUESE  IN  CHEMISTRY 

some  extent  in  every  liquid,  but  in  many  cases  the  solubility 
is  so  small  that  it  cannot  be  measured.  In  other  cases  it 
is  very  great.  Some  solids  dissolve  in  less  than  their  own 
weight  of  water,  and  some  gases,  such  as  ammonia,  dissolve 
to  the  extent  of  1000  volumes  in  1  volume  of  water. 

2.  Nature  of  the  solvent.    The  nature  of  the  solvent  is  no 
less  important.    Water,  alcohol,  and  ether  each  have  their 
own  peculiar  solvent  power.    Water  is  probably  the  most 
general  solvent  for  all  classes  of  materials,  and  alcohol  is 
perhaps  next  to  it.    Ether,   chloroform,   and  benzene  are 
good  solvents  for  organic  substances  such  as  fats,  waxes, 
and  oils. 

3.  Temperature.    The  weight  of  a  solid  which  a  given 
liquid  can  dissolve  varies  with  the  temperature.    Usually 
it  increases  rapidly  as  the  temperature  rises,  so  that  the 
boiling  liquid  dissolves  several  times  the  weight  which  the 
cold  liquid  will  dissolve.    In  some  instances,  as  in  the  case 
of  the  solubility  of  common  salt  in  water,  the  temperature 
has  little  influence,  and  a  few  solids  are  more  soluble  in  cold 
water  than  in  hot. 

In  the  case  of  gases,  on  the  other  hand,  the  lower  the 
temperature  of  the  liquid  the  larger  the  quantity  of  gas 
which  it  can  dissolve.  At  0°,  1000  volumes  of  water  will 
dissolve  41.14  volumes  of  oxygen;  at  50°,  18.37  volumes; 
at  100°,  none  at  all.  While  most  gases  can  be  expelled 
from  a  liquid  by  boiling  the  solution,  some  cannot.  For 
example,  it  is  not  possible  by  boiling  to  expel  hydrogen 
chloride  gas  completely  from  its  solution. 

Tables  of  solubilities.  For  convenience  of  reference  the  facts 
known  about  the  solubilities  of  various  substances  have  been 
collected  into  tables  of  solubilities,  and  these  are  constantly  used 
by  the  chemist.  Tables  giving  the  solubilities  of  a  few  of  the 
most  familiar  substances  will  be  found  in  the  Appendix. 


SOLUTIONS  AND  IONIZATION 


99 


4.  Pressure.  Change  of  pressure  has  little  effect  upon  the 
solubility  of  a  solid,  but  greatly  influences  that  of  a  gas. 
The  weight  of  a  gas  which  dissolves  in  a  given  case  is  pro- 
portional to  the  pressure  exerted  upon  the  gas  (Henry's  law). 
If  the  pressure  is  doubled,  the  weight  of  the  gas  going  into 
solution  is  doubled ;  if  the  pressure  is  diminished  one  half, 
then  but  half  as  much  gas  will  dissolve.  Under  high  pressure 
large  quantities  of  a  gas  can  be  dissolved  in  a  liquid,  and 
when  the  pressure  is  removed, 
the  escape  of  the  gas  causes  the 
liquid  to  foam,  or  effervesce,  as 
in  the  familiar  example  of  soda 
water. 

Colloidal  suspensions.  In  some 
cases  a  solid  may  be  distributed 
through  a  liquid  in  so  fine  a 
state  of  division  that  it  does 
not  render  the  liquid  noticeably 
cloudy,  and  does  not  settle  out 
even  after  long  standing ;  yet  it 
does  not  form  a  true  solution. 
This  is  called  a  colloidal  suspension,  or  colloidal  solution,  and 
the  solid  forming  it  is  called  a  colloid.  Starch,  gums,  resins, 
gelatin,  and  glue  are  typical  colloids. 

That  the  material  in  a  colloidal  suspension  is  not  really 
in  solution  is  shown  by  the  fact  that  a  strong  beam  of  light 
directed  through  such  a  suspension  has  a  bright  path,  like 
a  sunbeam  in  a  dark,  dusty  room,  while  its  path  through  a 
true  solution  is  invisible.  Fig.  49  is  a  photograph  of  a  beam 
of  light  from  a  lantern,  shining  through  a  colloidal  solution. 

Emulsions.  When  two  liquids  which  do  not  freely  mix 
with  each  other,  such  as  water  and  oil,  are  shaken  violently 
together,  a  milky  liquid  results,  consisting  of  very  small 


FIG.  49.  A  beam  of  light  shin- 
ing through  a  colloidal  solution 


100 


FIRST  COURSE  IN  CHEMISTRY 


droplets  of  the  one  liquid  suspended  in  the  other.  As  a 
rule  the  milkiness  quickly  disappears,  the  lighter  liquid 
collecting  as  a  layer  on  top  of  the  heavier  (Fig.  50,  A). 
If  a  colloidal  material,  such  as  soap,  is  added  before  shaking, 

this  separation  is  very  slow,  and 
the  milky  liquid  is  called  an 
emulsion  (Fig.  50,  IT).  Milk  is 
a  typical  emulsion,  the  butter 
fat  being  emulsified  by  the  col- 
loidal materials  present  in  the 
milk. 

Characteristic  properties  of 
solutions.  A  few  general  state- 
ments may  be  made  in  reference 
to  some  characteristic  proper- 
ties of  solutions. 

1.  Distribution  of  the  solid  in 
the  liquid.  A  solid,  when  dis- 
solved, tends  to  distribute  itself 
uniformly  through  the  liquid, 
so  that  every  part  of  the  solu- 
tion has  the  same  concentration. 
The  process  goes  on  very  slowly 
unless  hastened  by  stirring  or 
shaking  the  solution.  If  a  few 
crystals  of  a  highly  colored 
substance,  such  as  potassium  permanganate,  are  placed  in 
the  bottom  of  a  tall  vessel  full  of  water,  it  will  take  weeks 
for  the  solution  to  become  uniformly  colored. 

2.  Boiling  point  of  solutions.  The  boiling  point  of  a  liquid 
is  raised  by  the  presence  of  a  substance  dissolved  in  it. 
In  general  the  extent  to  which  the  boiling  point  of  a 
solvent  is  raised  by  a  given  substance  is  proportional  to 


B 

Emulsions 

In  A  the  oil  floats  on  the  water. 

If  soap  is  added  and  the  mixture 

shaken,  an  emulsion  is  formed  as 

shown  in  B 


SOLUTIONS  AND 

the  molecular  concentration  of  the  solution ;  that  is,  to 
the  number  of  gram-molecular  weights  of  the  substance 
dissolved  in  a  definite  weight  of  the  solvent. 

3.  Freezing  point  of  solutions.  The  freezing  point  of  a 
liquid  is  lowered  by  the  presence  of  a  substance  dissolved 
in  it.  The  lowering  of  the  freezing  point  obeys  a  law 
similar  to  the  one  which  holds  for  the  raising  of  the  boil- 
ing point,  the  extent  of  the  lowering  being  proportional  to 
the  molecular  concentration  of  the  solution. 

Electrolysis  of  solutions.  Pure  water  does  not  appreci- 
ably conduct  the  electric  current.  If,  however,  certain 
substances,  such  as  common  salt  or  acids,  are  dissolved  in 
the  water,  the  resulting  solutions 
are  found  to  be  good  conductors 
and  are  called  electrolytes.  When 
the  current  passes  through  an 
electrolyte,  some  chemical  change 
always  takes  place.  This  change 

is  called  electrolysis.  FlG-  61;  The  Process  of 

9  electrolysis 

The  general  method  used  in 

the  electrolysis  of  a  solution  is  illustrated,  in  Fig.  51.  Two 
plates  or  rods,  A  and  B,  made  of  suitable  material,  are  con- 
nected with  the  wires  from  a  battery  (or  dynamo)  C  and 
dipped  into  the  electrolyte,  as  shown  in  the  figure.  These 
plates  or  rods  are  called  electrodes.  The  electrode  B  con- 
nected with  the  negative  pole  of  the  battery  is  the  negative 
electrode,  or  cathode,  while  that  connected  with  the  positive 
pole  A  is  the  positive  electrode,  or  anode. 

Theory  of  ionization.  The  facts  discovered  in  connection 
with  electrolysis,  together  with  many  others,  have  led 
chemists  to  adopt  a  theory  of  solutions  called  the  theory  of 
ionization.  This  theory  was  first  proposed  by  the  Swedish 
chemist  Arrhenius  (Fig.  52).  Its  main  points  are  as  follows: 


102 


ABLEST  CD'UBSE  IN  CHEMISTRY 


1.  Formation  of  ions.   The  molecules  of  many  compounds, 
when  dissolved  in  water,  fall  apart,  or  dissociate,  into  two 
or  more  parts,  called  ions.    Thus,  sodium  nitrate  (NaNO3) 
dissociates  into  the  ions  Na  and  NOg ;    sodium  chloride 
(Nad),  into  the  ions  Na  and  Cl.    These  ions  move  about 
in  the  solution  independently  of  each  other  like  independ- 
ent molecules,  and  for  this 
reason  were  given  the  name 
ion,  which  means  "  wanderer." 
2.   The  electrical  charge  of 
ions.     An   ion   differs   from 
an  atom  or  molecule  in  that 
it  carries  a  large  electrical 
charge.     It  is   evident   that 
the  sodium  ion  must  differ 
in  some  important  way  from 
ordinary  sodium,  for  sodium 
ions,  formed  from  ordinary 
salt,  give  no  visible  evidence 
of  their  presence  in  water, 
whereas  metallic  sodium  at 
once  decomposes  the  water. 
The  electrical  charge,  there- 
fore,   greatly    modifies    the 
usual  chemical  properties  of 
the  element. 

3.  The  positive  charges  equal  the  negative  charges.  The  ions 
formed  by  the  dissociation  of  any  molecule  are  of  two  kinds: 
one  is  charged  with  positive  electricity  and  the  other  with 
negative.  The  sum  of  all  the  positive  charges  is  always 
equal  to  the  sum  of  all  the  negative  charges,  and  the  solu- 
tion as  a  whole  is  therefore  electrically  neutral.  If  we 
represent  ionization  by  the  usual  chemical  equations,  with 


FIG.  52.  Svante  August  Arrhenius 
(1859-) 

A  Swedish  chemist,  who  suggested 
the  theory  of  ionization 


SOLUTIONS  AND  IONIZATION 


103 


the  electrical  charges  indicated  by  plus  (+)  and  minus  (— ) 
signs  following  the  symbols,  the  ionization  of  sodium 
chloride  molecules  is  represented  thus : 


NaCl 


Cl- 


Those  ions  that  are  positively  charged  are  known  as  cations, 

while  those  that  are  negatively  charged  are  termed  anions. 

4.  Not  all  compounds  ionize.   It  is  assumed  that  only  those 

compounds  ionize  whose  solutions  are  electrolytes.    Thus, 


FIG.  53.  Laboratory  method  for  showing  whether  or  not  a  solution 
is  a  conductor 

salt  ionizes  when  dissolved  in  water,  for  it  has  been  found 
that  the  resulting  solution  is  a  very  good  electrolyte. 
Sugar,  on  the  other  hand,  does  not  ionize,  and  its  solution 
is  not  a  conductor  of  the  electric  current. 

Fig.  53  illustrates  a  very  convenient  apparatus  for  deter- 
mining whether  a  solution  is  a  good  conductor.  The  solution 
is  placed  in  the  bottle  A  and  the  electrodes  are  dipped  into  it. 
Connection  with  the  lighting  circuit  is  made  by  the  cord  and 
plug  B.  If  the  solution  is  a  good  conductor,  the  current  will 
flow  through  the  lamp  C,  which  will  then  glow. 


104  FIRST  COURSE  IN  CHEMISTRY 

The  theory  of  ionization  and  the  properties  of  solutions. 
In  order  to  be  of  value,  this  theory  must  be  in  accord  with 
the  chief  properties  of  solutions.  Let  us  now  see  if  the 
theory  is  in  harmony  with  certain  of  these  properties. 

The  theory  of  ionization  and  the  boiling  and  freezing  points 
of  solutions.  We  have  seen  that  the  boiling  point  of.  a  solu- 
tion of  a  substance  is  raised  in  proportion  to  the  number  of 
molecules  of  the  solute  present  in  the  solution. 

It  has  been  found,  however,  that  in  the  case  of  electrolytes 
the  boiling  point  is  raised  more  than  it  should  be  to  con- 
form to  this  law.  If  the  solute  dissociates  into  ions,  the 

reason   for   this   becomes 
clear.    Each  ion   has  the 
same  effect  on  the  boiling 
point  that  a  molecule  has, 
and  since  their  number  is 
always  greater   than   the 
FIG.  54.    The  electrolysis  of  sodium      number  of  molecules  from 
chloride 

which  they  were  formed, 

the  effect  on  the  boiling  point  is  abnormally  great. 

In  a  similar  way  the  theory  furnishes  an  explanation  of 
the  abnormal  lowering  of  the  freezing  point  of  electrolytes. 

The  theory  of  ionization  and  electrolysis.  The  changes 
taking  place  during  electrolysis  harmonize  very  completely 
with  the  theory  of  ionization.  This  will  become  clear  from 
a  study  of  the  following  examples : 

1.  Electrolysis  of  sodium  chloride.  Fig.  54  represents  a 
vessel  in  which  the  electrolyte  is  a  solution  of  sodium 
chloride  (NaCl).  According  to  the  theory  of  ionization, 
the  molecules  of  sodium  chloride  dissociate  into  the  ions 
Na+  and  Cl~.  Since  the  cathode  B  has  a  large  negative 
charge  derived  from  the  battery  (7,  the  Na+  ions  are  attracted 
to  it.  On  coming  in  contact  with  the  cathode,  they  give  up 


SOLUTIONS  AND  IONIZATION  105 

their  positive  charge  and  are  then  ordinary  sodium  atoms. 
They  immediately  decompose  the  water  according  to  the 
equation  2Na+ 2H>O  _^  2  NaOH +  Hf 

and  hydrogen  gas  is  evolved  from  the  surface  of  the  cathode. 
In  a  similar  way  the  chlorine  ions  (Cl~)  are  attracted  to 
the  positively  charged  anode  A,  and  upon  giving  up  their 
charge  to  it  they  are  set  free  as  chlorine  atoms  and  may 
either  combine  with  each  other  to  form  molecules  of  chlorine 
gas,  or  may  attack  the  water  as  represented  in  the  equation 

4  Cl  +  2  H2O  — *  4  HC1  +  O2 

It  is  to  be  carefully  noted  that  the  current  does  not  bring 
about  the  decomposition  of  the  solute  into  ions,  but  that  it  can 
pass  through  the  solution  only  when  ions  are  already  present. 

2.  Electrolysis  of  water.  The  reason  for  the  addition  of 
sulfuric  acid  to  water  in  the  preparation  of  oxygen  and 
hydrogen  by  electrolysis  (p.  17)  can  now  be  made  clear. 
Water  itself  is  not  an  electrolyte  to  an  appreciable  extent, 
for  it  does  not  form  enough  ions  to  carry  a  current.  Sulfuric 
acid  (H2SO4)  dissolved  in  water  is  an  electrolyte,  and  dis- 
sociates into  the  ions  2  H+  and  SO4~~,  each  SO4~~  ion  hav- 
ing two  negative  charges.  In  the  process  of  electrolysis  of 
the  solution  the  H+  ions  travel  to  the  cathode,  and  on  being 
discharged,  escape  as  hydrogen  gas.  The  SO4~~  ions,  when 
discharged  at  the  anode,  act  upon  the  water,  setting  free 
oxygen  and  once  more  forming  sulfuric  acid: 

2  S04  +  2  H20  — >•  2  H2S04  +  O2 

The  sulfuric  acid  can  again  ionize  and  the  process  repeat 
itself  as  long  as  any  water  is  left.  Hence  the  hydrogen  and 
oxygen  set  free  in  the  electrolysis  of  water  really  come 
directly  from  the  acid  but  indirectly  from  the  water. 


106  FIRST  COURSE  IN  CHEMISTRY 

Properties  of  electrolytes  dependent  upon  the  ions  present. 
When  a  substance  capable  of  forming  ions  is  dissolved  in 
water,  the  properties  of  the  solution  will  depend  upon  two 
factors :  (1)  the  ions  formed  from  the  substance ;  (2)  the 
undissociated  molecules.  Since  the  ions  are  usually  more 
active  chemically  .than  the  molecules,  most  of  the  chemical 
properties  of  an  electrolyte  are  due  to  the  ions  rather  than 
to  the  molecules. 

The  solutions  of  any  two  substances  which  give  the  same 
ion  will  have  certain  properties  in  common.  Thus,  all  solu- 
tions containing  the  copper  ion  Cu++  are  blue,  unless  the 
color  is  modified  by  the  presence  of  ions  or  molecules  having 
some  other  color. 

EXERCISES 

1.  Why  does  the  water  from  some  natural  springs  effervesce? 

2.  Why  does  not  the  water  of  the  ocean  freeze? 

3.  Why  does  shaking  or   stirring  make  a  solid   dissolve  more 
rapidly  in  a  liquid  ? 

4.  Why  will  vegetables  cook  faster  when  boiled  in  strong  salt 
water  than  when  boiled  in  soft  water? 

5.  How  do  you  explain  the  foaming  of  soda  water? 

6.  Account  for  the    fact  that,  sugar   sometimes   deposits  from 
sirups,  even  when  no  evaporation  has  taken  place. 

7.  How  did  the  ocean  become  salty  ? 

8.  What  is  the  solubility  of  CO2  in  water  (see  Appendix)  ?  On 
drawing  a  liter  of  soda  water  from  a  fountain,  500  cc.  of  CO2  escaped. 
Assuming  that  the  temperature  was  0°,  what  was  the  pressure  upon 
the  gas  in  the  fountain?    .4ns.  1.29  atmospheres. 

9.  10  g.  of  common  salt  was  dissolved  in  water  and  the  solu- 
tion evaporated  to  dryness ;  what  weight  of  solid  was  left  ?  Am.  10  g. 
10  g.  of  zinc  was  dissolved  in  hydrochloric  acid  and  the  solution 
evaporated  to  dryness;  what  weight  of  solid  was  left?   Ans.  20.8  g. 

TOPIC  FOR  THEMES 

Describe  the  details  of  a  soda-water  fountain  (from  examination). 


CHAPTER  XIII 
ACIDS,  BASES,  AND  SALTS;  NEUTRALIZATION 

Acids,  bases,  and  salts.  The  three  classes  of  compounds 
known  as  acids,  bases,  and  salts  include  the  great  majority 
of  the  compounds  with  which  we  shall  be  concerned.  A 
few  representatives  of  each  class  will  be  described  in  this 
chapter,  so  that  their  characteristic  properties  and  reactions 
may  be  made  clear. 

The  familiar  acids.  The  liquids  called  acids,  used  so 
largely  in  the  industries  and  in  chemical  laboratories,  are 
all  solutions  of  definite  compounds  in  water.  Hydrochloric 
acid  is  a  solution  of  a  gas  called  hydrogen  chloride,  which 
has  the  composition  expressed  in  the  formula  HC1.  Nitric 
acid  is  a  solution  of  the  liquid  known  as  hydrogen  nitrate, 
the  formula  of  which  is  HNOg.  Sulfuric  acid  is  a  solu- 
tion of  the  thick  oily  liquid  called  hydrogen  sulfate,  whose 
formula  is  Hj$O4.  For  most  purposes  it  is  not  neces- 
sary to  make  a  distinction  between  the  name  of  the  com- 
pound and  its  solution  in  water,  and  both  are  frequently 
called  acids. 

Characteristics  of  acids.  (1)  All  compounds  forming 
acids  in  solution  contain  hydrogen.  (2)  The  solutions  have 
a  sour  taste.  (3)  They  change  the  color  of  certain  sub- 
stances called  indicators.  Thus  blue  litmus  (a  dyestuff  ob- 
tained from  certain  lichens)  is  turned  red,  as  are  the  blue 
colors  of  most  flowers,  such  as  violets  and  corn  flowers. 
(4)  When  brought  in  contact  with  certain  metals,  acids 

107 


108  FIRST  COURSE  IN  CHEMISTRY 

evolve  hydrogen,  and  the  metal  dissolves.   With  hydrochlo- 
ric acid  and  zinc  the  reaction  is  represented  by  the  equation 

2  HC1  +  Zn >-  ZnCl2  +  H2 

lonization  of  acids.  When  dissolved  in  water,  the  com- 
pounds forming  acids  all  ionize  into  two  kinds  of  ions.  Of 
these  hydrogen  is  always  the  cation  (+),  while  the  remain- 
der of  the  molecule  is  the  anion  (— ),  thus : 

HC1 — ^H+  +  C1- 

Since  all  acids  produce  hydrogen  cations,  while  the 
anions  of  each  are  different,  the  properties  which  all  acids 
have  in  common  when  in  solution,  for  example,  taste,  and 
action  on  indicators  and  metals,  must  be  attributed  to  the 
hydrogen  ions.  We  may  therefore  define  an  acid  as  a  substance 
which  produces  hydrogen  ions  when  dissolved  in  water. 

Undissociated  acids.  When  compounds  that  form  acids  are 
perfectly  free  from  water,  or  are  dissolved  in  liquids,  like  ben- 
zene, which  do  not  have  the  power  of  dissociating  them  into 
ions,  they  have  no  real  acid  properties.  Under  these  circum- 
stances they^  do  not  affect  the  color  of  indicators  or  have  any 
of  the  properties  characteristic  of  acids. 

The  familiar  bases.  The  bases  most  used  in  the  laboratory 
are  sodium  hydroxide  (NaOH),  potassium  hydroxide  (KOH), 
and  calcium  hydroxide  (Ca(OH)2).  These  are  white  solids, 
soluble  in  water  —  the  latter  sparingly  so.  The  very  soluble 
bases  with  most  pronounced  basic  properties  (including  the 
three  just  mentioned)  are  sometimes  called  the  alkalies. 

Characteristics  of  bases.  (1)  All  compounds  forming 
bases  in  solution  contain  hydrogen  and  oxygen.  (2)  The 
solution  of  a  base  has  a  soapy  feel  and  a  bitter  taste. 
(3)  A  base  reverses  the  color  change  produced  in  indica- 
tors by  acids,  turning  red  litmus  blue. 


ACIDS,  BASES,  AND  SALTS  109 

lonization  of  bases.  When  dissolved  in  water,  the  mole- 
cules of  the  base  dissociate  into  two  kinds  of  ions.  One  of 
these  is  always  composed  of  the  group  O0  and  is  the  anion. 
It  is  called  the  hydroxyl  ion.  The  remainder  of  the  mole- 
cule, which  usually  consists  of  a  single  atom,  is  the  cation, 

thus:  NaOH — ^Na+,OH- 

Since  all  bases  produce  hydroxyl  anions,  while  the  cat- 
ions of  each  are  different,  the  properties  which  all  bases 
have  in  common  when  in  solution  must  be  due  to  the  hy- 
droxyl ions.  We  may  therefore  define  a  base  as  a  substance 
^hichj}roduces  hydroxyl  ions  when  dissolved  in  water. 

Neutralization.  When  an  acid  and  a  base  are  brought 
together  in  solution  in  proper  proportion,  the  characteristic 
properties  of  each  disappear.  As  a  rule  the  solution  tastes 
neither  sour  nor  bitter,  but  salty ;  it  has  no  effect  upon 
indicators.  There  can  therefore  be  neither  hydrogen  ions 
nor  hydroxyl  ions  in  the  solution.  This  action  of  an  acid 
on  a  base  is  called  neutralization. 

A  study  of  reactions  of  this  kind  has  shown  that  hydro- 
gen ions  and  hydroxyl  ions  cannot  exist  toge^ier  in  solu- 
tion to  any  appreciable  extent,  but  at  once  combine  to  form 
water.  The  following  equations  express  the  neutralization 
of  three  acids  by  three  bases,  water  being  formed  in  each 

Case :          Na+,  OH-  +  H+,  Cl~    ->-  Na+,  Cl~  +  H2O 
K+,  OH-  +  H+,  N03-  -  ->  K+,  N03-  +  H2O 
Ca+  +,  (OH-)2  +  (H+)2,  S04-  -     -+  Ca+  +,  SO4~  ~  +  2  H2O 

Neutralization  consists  in  the  union  of  the  hydrogen  ion  of  ^ 
(ft/  acid  with  the  hydroxyl  ion  of  a  base,  to  form  water. 

Salts.  It  will  be  noticed  that  in  neutralization  the  anion 
of  the  acid  and  the  cation  of  the  base  are  not  changed,  but 
remain  as  ions  in  the  solution.  If,  however,  the  water  is 


110  FIRST  COUESE  IN  CHEMISTRY 

expelled  by  evaporation,  these  two  ions  slowly  unite,  and 
when  the  water  becomes  saturated  with  the  substance  so 
produced,  it  begins  to  separate  in  the  form  of  a  solid  called 
a  salt.  Asalt,  therefore,  is  a  substance  formed  by  tlieunion  of 
the  anion^qf  an  add  with  the  cation  of  a  base. 

A  salt  may  also  be  obtained  by  the  action  of  a  metal  011 


and  we  may  say  that  a  salt  is  a  compound  obtained  by 
displacing  the  hydrogen  of  an  acid  by  a  metal.  The  salt 
derived  in  this  way  from  any  given  acid  is  said  to  be  a 
salt  of  that  acid.  Thus,  ZnSO4  is  a  salt  of  sulfuric  acid. 

Characteristics  of  salts.  (1)  From  the  definition  of  a 
salt  it  will  be  seen  that  there  is  no  element  or  group  of  ele- 
ments which  characterize  salts.  (2)  Salts  as  a  class  have  no 
peculiar  taste.  (3)  With  a  few  exceptions  (to  be  explained 
later)  they  are  without  action  on  indicators.  (4)  When 
dissolved  in  water  they  form  two  kinds  of  ions. 

Illustration  of  neutralization.  Sometimes  a  soil  becomes 
sour,  or  acid,  owing  to  the  formation  of  acids  which  are  often 
derived  from  decomposing  vegetable  matter.  Certain  plants, 
such  as  mosses  and  huckleberries,  will  thrive  in  acid  soil, 
but  grass,  clover,  and  grain  crops  will  not.  In  such  cases 
the  soil  must  be  sweetened  by  spreading  calcium  hydroxide 
(slaked  lime)  upon  it  to  neutralize  the  acids  present,  the 
process  being  called  liming  the  soil.  An  acid  soil  may  be 
detected  by  moistening  strips  of  blue  litmus  and  covering 
them  for  a  few  minutes  with  the  moist  soil.  Acids  will  turn 
the  blue  litmus  bright  red. 

Extent  of  ionization.  The  question  will  naturally  arise, 
When  an  acid,  base,  or  salt  dissolves  in  water,  do  all  the  mole- 
cules ionize,  or  only  some  of  them  ?  The  experiments  by  which 
this  question  is  answered  cannot  be  described  here.  It  has 


ACIDS,  BASES,  AND  SALTS  111 

been  found,  however,  that  only  a  fraction  of  the  molecules 
ionize.  The  percentage  which  ionizes  in  a  given  case  depends 
upon  several  conditions,  the  chief  of  which  are  as  follows : 

1.  The  concentration  of  the  solution.    In  concentrated  solu- 
tions only  a  very  small  percentage  of  the  molecules  ionize. 
As  the  solution  is  diluted  the  percentage  increases,  and  in 
dilute  solutions  it  may  be  very  large,  though  it  is  never 
complete  in  any  ordinary  solution. 

2.  The  nature  of  the  dissolved  compound.   At  equal  concen- 
trations substances  differ  much  among  themselves  in  per- 
centage of  ionization.   Most  salts  are  about  equally  ionized. 
Acids  and  bases,  on  the  contrary,  show  great  differences. 
Some  are  freely  ionized,  while  others  are  ionized  to  but  a 
slight  extent. 

Strength  of  acids  and  bases.  Since  acid  and  basic  proper- 
ties are  due  to  hydrogen  ions  and  hydroxyl  ions,  the  acid 
or  base  which  will  produce  the  greatest  percentage  of  these 
ions  at  a  given  concentration  must  be  regarded  as  the 
strongest  representative  of  its  class.  The  acids  and  bases 
described  in  the  foregoing  paragraphs  are  all  quite  strong. 
In  10  per  cent  solutions  about  half  of  the  molecules  are  dis- 
sociated into  ions,  and  this  is  also  approximately  the  extent 
to  which  most  salts  are  ionized  at  this  same  concentration. 

Methods  of  expressing  reactions  between  compounds  in 
solution.  Chemical  equations  representing  reactions  be- 
tween substances  in  solution  may  represent  the  details  of 
the  reaction,  or  they  may  simply  indicate  the  final  products 
formed.  Thus,  if  we  wish  to  call  attention  to  the  details  of 
the  reaction  between  sodium  hydroxide  and  hydrochloric 
acid  in  solution,  representing  the  ions  which  take  part  in 
the  reaction,  we  may  write  the  equation  as  follows : 

Na+,  OH-  -}-  H+,  Cl- >•  Na+,  Cl~  +  H2O 


112  FIKST  COURSE  IN  CHEMISTRY 

If  we  wish  simply  to  represent  substances  taking  part  in  the 
reaction  and  the  final  products  formed,  we  write  the  equation 

thus  :  NaOH  +  HC1  -  *•  NaCl  +  H2O 

Radicals.  It  has  been  emphasized  that  the  hydroxyl 
group  OH  always  forms  the  anion  of  a  base.  Similarly, 
the  group  NO3  forms  the  anion  of  nitric  acid  ;  the  group 
SO4,  the  anion  of  sulfuric  acid.  A  group  of  elements  which 
act  together  as  a  unit  in  chemical  action  is  called  a  radical. 

Some  of  these  radicals  have  been  given  special  names. 
Thus,  we  have  the  hydroxyl  radical  OH,  the  nitrate  radical 
NOQ,  and  the  sulfate  radical  SO.. 

3  4 

Displacement  series.  Upon  bringing  a  piece  of  zinc  into 
a  solution  of  an  acid,  zinc  passes  into  solution  and  hydrogen 
is  evolved  : 


111  like  manner,  when  zinc  is  placed  in  a  solution  of  a  salt 
of  copper,  such  as  the  sulfate  CuSO4,  zinc  passes  into  solu- 
tion, and  a  corresponding  quantity  of  copper  is  precipitated: 

Zn  +  CuS04  -  >•  ZnS04  +  Cu 

On  the  other  hand,  copper  has  no  effect  upon  a  solution  of 
zinc  sulfate. 

It  has  been  found  to  be  possible  to  arrange  hydrogen 
and  the  metals  in  a  table  in  such  a  way  that  any  ele- 
ment in  the  list  will  displace  any  one  below  it  from  its 
salts  and  will  in  turn  be  displaced  from  its  salts  by  any  one 
above  it.  This  list  is  called  the  displacement  series. 

DISPLACEMENT  SERIES 


1.  Potassium 
2.  Sodium 
3.  Magnesium 
4.  Aluminium 

5.  Zinc 
G.  Iron 
7.  Tin 
8.  Lead 

9.  Hydrogen 
10.  Copper 
11.  Antimony 
12.  Bismuth 

13.  Mercury 
14.  Silver 
15.  Platinum 
16.  Gold 

ACIDS,  BASES,  AND  SALTS  113 

This  table  enables  us  to  foretell  many  reactions.  For 
example,  from  the  positions  of  the  two  metals  we  should 
expect  magnesium  to  displace  tin  from  its  salts: 


We  should  not,  however,  expect  iron  to  displace  aluminium. 

It  is  of  especial  interest  to  notice  the  position  of  hy- 
drogen in  the  series.  All  the  metals  above  it  will  evolve 
hydrogen  from  solutions  of  its  salts  (acids),  while  those 
below  it  will  not.  In  the  latter  case,  if  any  action  takes 
place  it  must  be  preceded  by  oxidation. 

Names  of  acids,  bases,  and  salts.  Since  acids,  bases,  and 
salts  are  so  intimately  related  to  one  another,  it  is  very  ad- 
vantageous to  give  names  to  the  three  classes  in  accordance 
with  some  fixed  system.  The  system  universally  adopted 
is  as  follows  : 

Naming  of  bases.  All  bases  are  called  hydroxides.  They 
are  distinguished  from  each  other  by  prefixing  the  name  of 
the  element  which  is  in  combination  with  the  hydroxyl 
group.  Examples:  sodium  hydroxide  (NaOH);  calcium 
hydroxide  (Ca(OH)2);  copper  hydroxide  (Cu(OH)2). 

Naming  of  acids.  The  method  of  naming  acids  depends 
upon  whether  the  acid  consists  of  two  elements  or  of  three. 

1.  Binary  acids.    Acids  containing  only  one  element  in 
addition  to  hydrogen   are  called  Unary  acids.     They  are 
given  names  consisting  of  the  prefix  hydro-,  the  name  of 
the  second  element  present,  and  the  termination  -ic.   Exam- 
ples:  hydrochloric  acid  (HC1);  hydrosulfuric  acid  (H2S). 

2.  Ternary  acids.    In  addition  to  the  two  elements  pres- 
ent  in    binary    acids,    most    acids    also    contain    oxygen. 
These  acids  therefore  consist  of  three  elements  and  are 
called  ternary  acids.     It  usually  happens   that   the   same 
three  elements  can  unite  in  different  proportions  to  make 


114 


FIKST  COURSE  IN  CHEMISTRY 


several  different  acids.  The  most  familiar  one  of  these  is 
given  a  name  ending  in  the  suffix  4c,  while  the  one  with 
less  oxygen  is  given  a  similar  name,  but  ending  in  the 
suffix  -ous.  Examples :  nitric  acid  (HNOg) ;  nitrous  acid 
(HN02). 

Naming  of  salts.  A  salt  derived  from  a  binary  acid  is 
given  a  name  consisting  of  the  names  of  the  two  elements 
composing  it,  with  the  termination  -ide.  Example  :  sodium 
chloride  (NaCl).  All  other  binary  compounds  are  named 
in  the  same  way. 

A  salt  of  a  ternary  acid  is  named  in  accordance  with  the 
acid  from  which  it  is  derived.  A  ternary  acid  with  the  ter- 
mination -ic  gives  a  salt  with  the  name  ending  in  -ate,  while 
an  acid  with  the  termination  -ous  gives  a  salt  with  the  name 
ending  in  -ite.  The  following  table  will  make  the  application 
of  these  principles  clear : 


ACID 

FORMULA 

SALT 

FORMULA 

Hydrochloric 

IIC1 

Sodium  chloride 

NaCl 

Chlorous 

HC1O2 

Sodium  chlorite 

NaClO2 

Chloric 

HC103 

Sodium  chlorate 

NaClO8 

EXERCISES 

1.  Name  three  edible  substances  that  taste  sour.   How  could  you 
prove  the  sour  taste  to  be  due  to  acids  ? 

2.  If  your  clothing  should  become  spotted  with  acids  in  the  labo- 
ratory, how  would  you  try  to  remove  the  spots  ? 

3.  Compose  a  definition  of  a  base,  different  from  the  one  given 
in  the  text ;  of  an  acid ;  of  a  salt. 

4.  What  weight  of  sodium  hydroxide  will  be  neutralized  by  50  g. 
of  hydrochloric  acid?    Ans.  54.85  g. 

5.  What  weight  of  sulfuric  acid  will  be  neutralized  by  40  g.  of 
sodium  hydroxide?    Ans.  49.03  g. 


ACIDS,  BASES,  AND  SALTS  115 

6.  Which  will  neutralize  the  most  hydrochloric  acid,  50  g.  of 
sodium  hydroxide  or  50  g.  of  calcium  hydroxide  ? 

7.  What  weight  of  copper  will  be  precipitated  from  a  solution  of 
copper  sulf ate  by  GO  g.  of  zinc  ?    A  ns.  58.34  g. 

8.  The  most  common  of  the  ternary  acids  of  sulfur  has  the  for- 
mula H2SO4.    What  is  its  name  ?    To  what  class  (acids,  bases,  salts) 
do  each  of  the  following  compounds  belong  :  H2SO3,  Na2SO4,  Na2SO8, 
Zn(OH)2,  A1(OH)3?    Give  the  name  of  each. 

9.  Account  for  the  fact  that  calcium  hydroxide,  when  perfectly  dry, 
has  no  action  on  litmus  paper. 

TOPIC  FOR  THEMES 

The  liming  of  soils.    (Write  to  the  director  of  your  experiment 
station  for  bulletins.) 


CHAPTER  XIV 
VALENCE 

Definition  of  valence.  A  comparison  of  the  composition 
of  the  compounds  of  hydrogen  with  the  other  elements 
brings  to  light  an  interesting  fact  illustrated  in  the  formulas 

HC1  H2O  H3N  H4C 

(hydrogen  chloride)  (water)  (ammonia)  (marsh  gas) 

It  will  be  seen  that  the  various  kinds  of  atoms  differ  in 
the  number  of  hydrogen  atoms  that  they  are  able  to  hold 
in  combination.  An  atom  of  chlorine  combines  with  but 
one  hydrogen  atom,  an  atom  of  oxygen  with  two,  one  of 
nitrogen  with  three,  and  one  of  carbon  with  four.  It  is 
convenient  to  have  a  name  to  designate  that  property  of 
an  element  that  determines  the  number  of  hydrogen  atoms 
that  its  atom  can  hold  in  combination.  It  is  called  the 
valence  of  an  element. 

Classifications  of  elements  according  to  their  valences. 
Valence  is  merely  the  ratio  between  the  numbers  of  two 
kinds  of  atoms  which  combine,  and  gives  no  information  in 
regard  to  the  intensity  of  affinity  between  the  atoms.  To 
express  the  valence  of  an  element,  we  must  select  some 
standard  for  comparison,  just  as  we  do  in  the  case  of  any 
other  numerical  quantity.  Since  one  atom  of  hydrogen 
never  combines  with  more  than  one  atom  of  any  other  ele- 
ment, hydrogen  is  selected  as  the  standard  and  is  said  to 
be  univalent.  Other  elements,  such  as  chlorine,  iodine,  and 
sodium,  which  combine  with  hydrogen  atom  for  atom  (HC1), 

116 


VALENCE  117 

are  likewise  said  to  be  univalent.  On  the  other  hand,  ele- 
ments such  as  oxygen,  sulfur,  and  calcium,  one  atom  of 
which  combines  with  two  atoms  of  hydrogen  or  of  other  uni- 
valent elements,  are  said  to  be  bivalent.  Similarly,  we  have 
trivalent  elements,  such  as  nitrogen,  and  quadrivalent  ones, 
such  as  carbon.  No  element  is  known  whose  valence  exceeds 
8,  and  with  the  majority  of  elements  it  does  not  exceed  4. 

How  the  valence  of  an  element  may  be  inferred.  The 
valence  of  an  element  may  readily  be  inferred  from  the 
formula  of  the  compound  which  it  forms  with  hydrogen  or 
other  elements  whose  valence  is  known.  The  same  method 
applies  to  radicals,  for  these  are  groups  of  atoms  which 
act  as  a  unit  and  on  this  account  may  be  assigned  valences 
just  as  though  they  were  elements.  Thus  the  radical  OH 
must  be  univalent,  since  it  combines  with  one  atom  of  the 
univalent  element  sodium,  as  shown  by  the  formula  NaOH  ; 
the  radical  SO4  must  be  bivalent,  since  it  combines  with  two 
atoms  of  hydrogen,  as  shown  by  the  formula  Hj$O4. 

Valence  and  combination  ratios.  Elements  or  radicals  hav- 
ing the  same  valence  combine  with  each  other  atom  for  atom. 
When  two  elements  having  different  valences  combine,  the 
combination  will  take  place  between  such  numbers  of 
atoms  as  have  equal  valences.  Thus  one  atom  of  a  bivalent 
element  or  radical,  such  as  zinc,  has  the  same  number  of 
valences  as  two  atoms  of  a  univalent  element,  such  as  chlo- 
rine, namely  two.  Accordingly  when  such  elements  unite, 
the  union  will  be  between  one  atom  of  the  bivalent  element 
and  two  of  the  univalent  element,  as  shown  by  the  formula 
ZnCl2.  When  a  trivalent  element  unites  with  a  bivalent 
element  or  radical,  the  union  will  be  between  two  atoms  of 
the  trivalent  element  (total  valence  of  6)  and  three  atoms 
of  the  bivalent  (total  valence  of  6),  as  expressed  in  such 
formulas  as  A12O3  and  A12(SO4)3. 


118  FIRST  COURSE  IN  CHEMISTRY 

It  is  evident  therefore  that,  knowing  the  valences  of  two 
elements  or  radicals,  we  can  tell  the  probable  formula  of 
the  compound  which  will  be  formed  by  their  union ;  con- 
versely, knowing  the  formula  of  the  compound  and  the 
valence  of  one  of  the  elements  present,  the  valence  of  the 
other  may  be  inferred.  For  example,  if  we  know  that  iron 
is  trivalent  in  the  compound  Fe2S3,  then  the  sulfur  must  be 
bivalent ;  for  the  three  atoms  of  sulfur  will  have  the  same 
total  valence  as  the  two  atoms  of  iron,  namely,  6.  One  atom 
of  sulfur  therefore  has  a  valence  of  2.  In  the  case  of  a 
few  compounds,  such  as  hydrogen  peroxide  (H2O2),  this 
principle  does  not  hold  good,  but  the  discussion  of  these 
is  beyond  the  scope  of  an  elementary  text. 

Variable  valence.  Since  many  of  the  elements  unite  with 
each  other  in  more  than  one  ratio,  it  follows  that  some  of 
the  elements  may  have  more  than  one  valence.  Thus,  in  the 
compound  CO  carbon  has  a  valence  of  2,  while  in  the  com- 
pound CO2  it  has  a  valence  of  4.  Similarly,  in  the  oxides 
FeO  and  Fe2O3  iron  has  a  valence  of  2  and  3  respectively. 
It  is  the  custom,  therefore,  to  speak  of  the  valence  of  an  ele- 
ment in  a  given  compound,  or  toward  some  given  element. 

Valence  and  the  replacing  power  of  atoms.  Just  as  ele- 
ments having  the  same  valence  combine  with  each  other 
atom  for  atom,  so,  if  they  replace  each  other  in  a  chemical 
reaction,  they  will  do  so  in  the  same  ratio.  Thus,  one  atom 
of  bivalent  zinc  displaces  two  atoms  of  univaleiit  hydrogen, 
as  is  shown  in  the  following  equations : 

Zn  +  H2SO4  -  ->-  ZnSO4  +  H2 
Zn(OH)2  +  H2S04 *•  ZnS04  +  2  H2O 

Similarly,  one  atom  of  bivalent  calcium  displaces  one  atom 
of  bivalent  zinc : 

CaCl2  +  ZnSO4 >•  CaSO4  +  ZnCl2 


VALENCE 


119 


Since  many  reactions,  like  those  above,  consist  in  an  inter- 
change of  two  elements,  it  is  evident  that  a  knowledge  of 
the  valence  of  the  elements  will  assist  us  in  writing  the 
equations  for  the  reactions. 

Valence  and  charges  on  ions.  We  have  represented  ions 
as  formed  by  the  dissociation  of  molecules  and  as  carrying 
electrical  charges,  thus : 


HC1 


H+  +  Cl- 


Since  hydrogen  and  chlorine  are  univalent,  it  will  be  seen 
that  their  valence  is  the  same  as  the  number  of  electrical 
charges  they  carry  as  ions.  Similarly,  calcium  sulfate  ionizes 
as  follows :  ~  ^^  ~  ++  Q~  __ 

In  this  case  both  the  calcium  ion  and  the  sulfate  ion  carry 
two  electrical  charges,  and  each  is  bivalent. 

Table  of  valences.  It  will  be  convenient  for  reference  to 
tabulate  the  valences  of  the  most  familiar  elements  and 
radicals.  In  the  table  which  follows,  some  elements  will  be 
found  in  several  different  columns,  since  they  have  more 
than  one  valence.  The  table  also  indicates  the  number  of 
electrical  charges  carried  when  the  element  or  radical  acts 


as  an  ion. 


TABLE  OF  VALENCES 


VALENCE 

POSITIVE  IONS 

NEGATIVE  IONS 

NOT  IONS 

1 

H,  Na,  K,  Ag,  NH4 

Cl,  Br,  I,  OH,  NO3 

2 

Ca,  Ba,  Mg,  Zn,  Hg, 

S,  S04,  C03 

O 

Cu,  Fe,  Sn 

3 

Al,  Bi,  Sb,  Fe 

.P°4 

N,  P 

4 

Sn 

SiO4 

C,  Si,  S 

5 

X,  P,  As,  Sb 

6 

S 

120  FIRST  COURSE  IN  CHEMISTRY 

EXERCISES 

1.  Write  the  formulas  of  the  compounds  that  sodium  (univalent) 
forms  with  each  of  the  following  elements  or  radicals :  Cl (univalent) ; 
S(bivalent);  SO4 (bivalent)  ;  PO4(trivalent). 

2.  In  the  compounds  whose  formulas  follow,  the  valence  of  one 
of  the  elements  is  indicated  by  a  figure  placed  over  the  symbol. 
What  is  the   valence   of   each  of    the   other  elements  or  radicals 
present  in  the  compounds? 

222  2  11 

MgBr2;  CaO;  Ca(OH)2;  Ba3(PO4)2;  KNO3;  H3P 

Verify  your  results  by  reference  to  the  table. 

3.  Complete  the  following  equations,  supposing  that  the  sodium 
or  the    calcium    in   the   first   compound  changes    places  with  the 
hydrogen  of  the  second : 

XaOH  +  HC1 >-- 

NaOH  +  H2SO4 y  - 

NaOH  +  H3PO4 ^- 

Ca(OH)2  +  HC1 >-  - 

Ca(OH)2  +  H2SO4 >-  - 

Ca(OH)2  +  H3PO4 >•  — 


CHAPTER  XV 
COMPOUNDS  OF  NITROGEN 

Occurrence.  The  large  quantity  of  nitrogen  occurring  in 
the  atmosphere  (p.  80)  is  practically  all  in  the  free  state. 
In  the  materials  composing  the  earth's  crust,  on  the  other 
hand,  there  occur  in  certain  localities  considerable  deposits  of 
nitrogen  compounds,  especially  of  sodium  nitrate  (NaNOg). 
The  latter  compound  is  found  in  large  quantities  in  Chile, 
and  serves  as  the  material  from  which  many  of  the  other 
compounds  of  nitrogen  are  prepared.  Small  quantities  of 
nitrogenous  compounds  also  occur  in  all  productive  soils. 
From  these  soils  the  nitrogen  is  taken  up  by  plants  and 
built  into  complex  compounds  known  as  protein  matter. 
Animals  feeding  on  these  plants  assimilate  the  nitrogenous 
matter,  which  becomes  an  essential  part  of  the  animal  tissue 
(p.  83). 

While  a  great  many  compounds  of  nitrogen  are  known, 
it  is  desirable  at  this  time  to  discuss  only  a  few  of  the 
simpler  ones. 

AMMONIA  (NH3) 

Several  compounds  consisting  exclusively  of  nitrogen 
and  hydrogen  are  known,  but  only  one,  ammonia,  need  be 
considered  here. 

Preparation  of  ammonia.  Ammonia  is  prepared  in  the 
laboratory  by  a  different  method  from  the  one  which  is 
used  commercially, 

121 


122  FIRST  COURSE  IN  CHEMISTRY 

1.  Laboratory  method.  In  the  laboratory,  ammonia  is  pre- 
pared from  ammonium  chloride  (NH4C1),  a  white  solid  ob- 
tained in  the  manufacture  of  coal  gas  (p.  210).  In  this 
compound  the  group  NH4  acts  as  a  univalent  radical  and  is 
known  as  ammonium.  When  ammonium  chloride  is  warmed 
with  sodium  hydroxide,  the  ammonium  radical  and  sodium 
change  places,  the  reaction  being  expressed  in  the  following 
equation : 

NH  Cl  +  NaOH *•  NaCl  +  NH  OH 


The  ammonium  hydroxide  (NH4OH)  so  formed  is  unstable 
and  breaks  down  into  water  and  ammonia  : 


NH4OH >-  NH8  +  H2O 

Calcium  hydroxide  (Ca(OH)2)  is  frequently  used  in  place  of 
the  more  expensive  sodium  hydroxide,  the  equations  being 

2  NH4C1  +  Ca(OH)2 >•  CaCl0  +  2  NH4OH 

2  NH4OH >•  2  II26  +  2  NH8 

The  ammonium  chloride  and  calcium  hydroxide  are  mixed 
together  and  placed  in  a  flask  A,  arranged  as  shown  in  Fig-.  55. 
The  mixture  is  gently  warmed,  when  ammonia  is  evolved  as 
a  gas  and,  being  much  lighter  than  air,  is  collected  in  B  by 
displacement  of  air,  as  shown  in  the  diagram. 

2.  Commercial  method.  Most  of  the  ammonia  of  commerce 
is  obtained  by  heating  soft  coal  in  the  absence  of  air,  as 
is  done  in  the  manufacture  of  illuminating  gas  or  coke 
(p.  210).  Among  the  by-products  of  this  operation  is  an 
impure  solution  of  ammonia  in  water,  known  as  gas 
liquor.  This  is  treated  with  lime  and  heated,  the  ammonia 
being  driven  out  as  a  gas.  The  gas  may  be  dissolved  again 
in  pure,  cold  water,  forming  aqua  ammonia,  or  the  ammonia 
water  of  commerce. 


COMPOUNDS  OF  NITROGEN 


123 


Preparation  from  hydrogen  and  nitrogen.  When  electric  sparks 
are  passed  for  some  time  through  a  mixture  of  hydrogen  and 
nitrogen,  a  small  percentage  of  the  two  elements  in  the  mix- 
ture is  changed  into  ammonia.  The  yield  of  ammonia  is  very 
small,  for  the  reason  that  the  compound  is  decomposed  by 
the  spark,  and  a  point  is  soon  reached  when  the  speed  of  its 
decomposition  is  as  great  as  the  speed  of  its  formation.  These 
facts  are  expressed  in 
the  equation 


FK 


This  recalls  the  similar 
relation  between  oxy- 
gen and  ozone. 

The  Haber  process. 
When  a  mixture  of  ni- 
trogen and  hydrogen 
subjected  to  a  pressure 
of  200  atmospheres  is 
heated  to  about  500°  in 
contact  with  finely  di- 
vided iron,  a  larger  per- 
centage of  ammonia  is 
formed.  This  is  dis- 
solved in  water  as  fast 

as  formed,  and  the  remaining  gases  are  again  conducted  over 
the  iron.  The  process  thus  becomes  continuous,  more  nitrogen 
and  hydrogen  being  introduced  as  needed.  Large  factories  are 
being  built  in  Germany  for  making  ammonia  by  this  process 
for  use  in  fertilizers.  The  nitrogen  used  is  obtained  from  the 
air,  and  the  hydrogen  from  water. 

Properties.  Under  ordinary  conditions  ammonia  is  a  gas 
which  is  little  more  than  half  as  heavy  as  air.  It  is  easily 
condensed  into  a  colorless  liquid,  and  can  now  be  purchased 
in  this  form.  The  gas  is  colorless  and  has  a  strong,  suffo- 
cating odor  often  noticed  about  decaying  organic  matter. 


55.    The  preparation  of  ammonia  in 
the  laboratory 


124 


FIRST  COURSE  IN  CHEMISTRY 


It  is  extremely  soluble  in  water,  1  liter  of  water  at  0°  and 
760  mm.  pressure  dissolving  1298  liters  of  the  gas,  and  at 
20°,  710  liters.  In  dissolving  such  large  volumes  of  the  gas 
the  water  expands  considerably,  so  that  the  density  of  the 
solution  is  less  than  that  of  water,  the  most  concentrated 
commercial  solutions  having  a  density  of  0.88. 

Chemical  conduct.    Ammonia  will  not  support  combus- 
tion,  nor  will  it  burn  under  ordinary  conditions.    In   an 
atmosphere  of  oxygen  it   burns  with  a  feeble,  yellowish 
n  flame.    When  quite 

dry,  it  is  not  a  very 

H  B  active  substance,  but 

(E  i 

I'-  ;•.'      •   ,  .   *    :•    •  .•    >•  ;•:.•';'•  J>  when  moist,  it  com- 

bines  with  a  great 
many  substances,  par- 
ticularly with  acids. 
Uses.  It  has  been 
stated  that  ammonia 
can  be  condensed  to 
a  liquid  by  the  appli- 
cation of  pressure. 
If  the  pressure  is  re- 
moved from  the  liquid  so  obtained,  this  liquid  rapidly 
passes  again  into  the  gaseous  state,  and  in  so  doing  absorbs 
a  great  deal  of  heat.  Advantage  is  taken  of  this  fact  in 
the  manufacture  of  ice.  Large  quantities  of  ammonia  are 
also  used  in  the  manufacture  of  ammonium  compounds. 

The  manufacture  of  ice.  Fig.  56  illustrates  the  method  of 
manufacturing  ice.  The  ammonia  gas  is  liquefied  in  the  pipes 
Aj  B  by  means  of  a  compression  pump.  The  heat  generated 
by  liquefaction  is  absorbed  by  water  flowing  over  the  pipes. 
The  pipes  lead  into  a  large  brine  tank,  a  cross  section  of 
which  is  shown  in  the  figure.  Into  the  brine  (concentrated 


FIG.  50.   Diagram  of  an  ammonia  ice  machine 


COMPOUNDS  OF  NITROGEN  125 

solution  of  calcium  chloride)  contained  in  this  tank  are  dipped 
the  vessels  D,  E,  F,  filled  with  pure  water.  The  pressure  is 
removed  from  the  liquid  ammonia  as  it  passes  through  the 
expansion  valve  C  into  the  pipes  immersed  in  the  brine,  and 
the  heat  absorbed  by  the  rapid  evaporation  of  the  liquid  lowers 
the  temperature  of  the  brine  below  zero.  The  water  in  D,  E,  F 


FIG.  57.    The  interior  of  a  cold-storage  room 

is  thereby  frozen  into  cakes  of  ice.  The  gaseous  ammonia  re- 
sulting from  the  evaporation  of  the  liquid  escapes  at  G  and 
is  again  liquefied,  and  the  process  is  repeated. 

Cold  storage.  The  temperature  of  a  room  may  be  kept  very 
low  by  surrounding  it  with  pipes  into  which  liquid  ammonia  is 
forced  and  allowed  to  vaporize.  The  vapor  is  conducted  away 
and  again  compressed  to  the  liquid  state  and  again  returned, 
so  that  the  process  becomes  continuous.  The  cold-storage 
plants  now  so  largely  used  for  preventing  the  decay  of  food 
products  are  constructed  on  this  principle.  Fig.  57  shows 


126  FIKST  COUESE  IN  CHEMISTRY 

a  corner  of  a  cold-storage  room.     The  pipes  containing  the 
ammonia  soon  become  covered  with  frozen  moisture. 

Ammonium  hydroxide  (NH4OH).  The  solution  of  am- 
monia in  water,  commonly  called  aqua  ammonia,  has  strong 
basic  properties.  It  turns  red  litmus  blue  ;  it  has  a  soapy 
feel  ;  it  neutralizes  acids,  forming  salts  with  them.  It  seems 
certain,  therefore,  that  when  ammonia  dissolves  in  water  it 
combines  chemically  with  it  according  to  the  equation 

H0—  >NHOH 


and  that  it  is  the  substance  NH4OH,  called  ammonium 
hydroxide,  which  has  the  basic  properties,  dissociating  into 
the  ions  NH4  and  OH.  The  separation  of  the  pure  hydrox- 
ide from  its  solutions  has  not  been  accomplished,  for  as 
the  solution  becomes  concentrated,  the  compound  decom- 
poses again  into  ammonia  and  water: 

NH  OH  -  >•  Nil  +  HO 

4  o  2 

The  solution  of  ammonia  in  water,  therefore,  constitutes 
a  state  of  equilibrium  between  ammonia  and  water,  on  the 
one  hand,  and  ammonium  hydroxide,  on  the  other.  This 
condition  is  conveniently  expressed  in  the  following  way: 


Aqua  ammonia  is  a  good  solvent  for  grease  and  is  a  famil- 
iar household  article. 

The  ammonium  radical.  The  univalent  radical  NH4  plays 
the  part  of  a  metal  in  many  chemical  reactions,  and  is  called 
ammonium.  The  ending  -ium  is  given  to  the  name  to  indicate 
the  metallic  properties  of  the  substance,  since  the  names  of  the 
metals  in  general  have  that  ending.  The  salts  formed  by  the 
action  of  the  base  ammonium  hydroxide  on  acids  are  called 


COMPOUNDS  OF  NITROGEN 


127 


ammonium  salts.    Thus,  with  hydrochloric  acid,  ammonium 
chloride  (NH4C1)  is  formed,  in  accordance  with  the  equation 

NH4OH  +  HC1  — *•  NH4C1  +  H2O 
Combination  of  nitrogen  with  hydrogen  by  volume.  Am- 
monia can  be  decomposed"  into  nitrogen  and  hydrogen  by 
passing  electric  sparks  through  the  gas.  Accurate  measure- 
ment has  shown  that  2  volumes  of  ammonia  yield  1  volume 
of  nitrogen  and  3  volumes  of  hydrogen.  These  relations  may 
be  represented  graphically  as  follows : 


NH8 

NH3 

H 


ACIDS  OF  NITROGEN 

The  most  important  of  the   acids   of  nitrogen   are  the 
following:  nitric  acid  (HNO3)  and  nitrous  acid  (HNO2). 

Nitricacid(HN03). 
This  acid  is  pre- 
pared from  sodium 
nitrate  (NaNO3), 
which  occurs  in  na- 
ture in  large  quan- 
tities. 

Preparation  of  ni- 
tric acid.  When  so- 
dium nitrate  is  treated 


FIG.  58.  The  preparation  of  nitric  acid  in 
the  laboratory 


with     concentrated, 

cold    sulfuric    acid, 

no    chemical    action 

seems  to  take  place.    If,  however,  the  mixture  is  heated 

in  a  retort  J,  nitric  acid  is  given  off  as  a  vapor  and  may  be 

easily  condensed  to  a  liquid  by  passing  the  vapor  into  a 

tube  B  surrounded  by  cold  water,  as  shown  in  Fig.  58. 


128 


FIKST  COURSE 


CHEMISTRY 


An  examination  of  the  liquid  left  in  the  retort  shows  that 
it  contains  sodium  acid  sulfate  (NaHSO4),  only  half  of  the 
hydrogen  of  sulfuric  acid  having  been  replaced  by  sodium. 
The  reaction  may  be  represented  by  the  equation 


NaNO. 


NaHSO4  +  UNO, 


The  commercial  preparation  of  nitric  acid.  Fig.  59  illustrates 
a  form  of  apparatus  used  in  the  preparation  of  nitric  acid  on 
a  large  scale.  Sodium  nitrate  and  sulfuric  acid  are  heated  in 

the  iron  retort. I. 
The  resulting  acid 
vapors  pass  in 
the  direction  indi- 
cated by  the  ar- 
rows, and  are 
condensed  in  the 
glass  tubes  B, 
which  are  covered 
with  cloth  kept 
cool  by  streams 
of  water.  These 
tubes  are  inclined 
so  that  the  liquid 
resulting  from  the 
condensation  of  the  vapors  runs  back  into  C  and  is  drawn  off 
into  the  large  vessel  D. 

Preparation  of  nitric  acid  from  air.  When  electric  dis- 
charge takes  place  through  a  mixture  of  oxygen  and  nitro- 
gen (air),  a  small  percentage  of  oxides  of  nitrogen  is  formed. 
This  can  be  increased  by  having  the  mixture  pass  through 
an  electric  arc  which  lias  been  drawn  out  to  a  great  size 
by  magnets  (Fig.  60).  The  oxides  so  obtained  combine  with 
water  to  form  dilute  nitric  acid.  This  method  for  preparing 
nitric  acid  (known  as  the  Birkeland  and  Eyde  process)  has 
come  into  extensive  use  in  recent  years  in  Norway,  since 


FIG.  59.  The  commercial  preparation  of  nitric  acid 


COMPOUNDS  OF  NITROGEN  129 

the  necessary  electrical  energy  is  generated  at  a  very  low 
cost  by  the  waterfalls  abounding  in  that  country.  The 
dilute  nitric  acid  obtained  is  neutralized  with  lime  (CaO) 
and  the  resulting  calcium  nitrate  sold  for  use  as  a  fertilizer 
under  the  name  air  saltpeter. 

Properties  of  nitric  acid.  Pure  nitric  acid  (hydrogen 
nitrate)  is  a  colorless  liquid  which  boils  at  about  86°  and 
has  a  density  of  1.56.  The  concentrated  acid  of  commerce 
contains  about  68  per  cent  of  the  acid,  the  remainder  being 
water.  Such  a  mixture  has  a  den- 
sity of  1.4.  The  concentrated  acid 
fumes  somewhat  in  moist  air,  and 
has  a  sharp,  choking  odor. 

Chemical  conduct.  The  most  im- 
portant chemical  reactions  of  nitric 
acid  are  the  following :  FIG.  60.  Form  of  the  elec- 

1.  Acid  action.    Nitric   acid    has     ^    ^rc    employed    in    the 

Birkeland  and  Eyde  process 
all  the  characteristics  of  a  strong 

acid.  It  changes  blue  litmus  red,  and  has  a  sour  taste  in 
dilute  solutions.  It  gives  hydrogen  ions  in  solution,  and 
neutralizes  bases,  forming  salts.  It  also  acts  upon  the  oxides 
of  most  metals,  forming  a  salt  and  water,  thus : 

CuO  +  2  HN03  — >-  Cu(N03)2  +  H2O 

2.  Decomposition  on  heating.    When  nitric  acid  is  boiled 
or  when  it  is  exposed  for  some  time  to  sunlight,  it  suffers 
a  partial  decomposition  according  to  the  equation 

4  HN03  — >-  2  HaO  +  4  NO2  +  O2 

The  substance  NO2  (called  nitrogen  dioxide)  is  a  brownish 
gas  which  is  readily  soluble  in  water  and  in  nitric  acid.  It 
therefore  dissolves  in  the  undecomposed  acid,  and  imparts 


130  FIRST  COURSE  IN  CHEMISTRY 

to  it  a  yellowish  or  reddish  color.  Concentrated  nitric  acid 
highly  charged  with  this  substance  is  colled  fuming  nitric  acid. 
3.  Oxidizing  action.  Because  of  its  easy  decomposition, 
nitric  acid  is  a  good  oxidizing  agent.  Under  ordinary 
circumstances,  when  acting  as  an  oxidizing  agent,  it  is 
decomposed  according  to  the  equation 

2  HN03  -  >•  H20  +  2  NO  +  3[O] 

The  oxygen  is  taken  up  by  the  substance  oxidized  and 
is  not  set  free,  which  fact  is  indicated  in  the  equation  by 
placing  the  symbol  for  oxygen  in  brackets.  Thus,  if  carbon 
is  oxidized  by  nitric  acid,  the  oxygen  combines  with  carbon, 
forming  carbon  dioxide  (CO0)  : 


C  +  2[0] 

4.  Action  on  metals.  Nitric  acid  acts  in  two  different  ways 
upon  metals:  (1)  With  dilute  nitric  acid  the  metals  above 
hydrogen  in  the  displacement  series  (p.  112),  such  as  zinc, 
evolve  hydrogen: 

Zn  +  2  HN03  -  *•  Zn(N03)2  +  H2 

(2)  With  concentrated  nitric  acid  none  of  the  metals  evolve 
hydrogen,  but  oxidation  takes  place  forming  a  nitrate  to- 
gether with  a  gas  called  nitric  oxide  (NO).  With  copper 
the  reaction  is  expressed  in  the  equation 

3  Cu  +  8  HNO3  -  *  3  Cu(NO8)2  +  2  NO  +  4  H2O 

A  similar  action  often  occurs  when  dilute  nitric  acid  acts 
upon  metals  below  hydrogen  in  the  displacement  series. 

Uses.  Nitric  acid  has  countless  uses  in  the  industries  and 
in  chemical  laboratories.  It  is  most  extensively  used  in  the 
manufacture  of  explosives  of  various  kinds  and  of  celluloid 
and  artificial  dyes. 


COMPOUNDS  OF  NITROGEN  131 

Salts  of  nitric  acid ;  nitrates.  The  salts  of  nitric  acid 
are  called  nitrates.  Many  of  these  salts  will  be  described 
in  the  study  of  the  metals.  They  are  all  soluble  in  water, 
and  when  heated  to  a  high  temperature,  undergo  decompo- 
sition. In  a  few  cases  a  nitrate,  on  being  heated,  evolves 
oxygen,  forming  a  nitrite: 

2  NaNOq  — >•  2  NaNO9  +  O9 

o  X  -    3 

In  other  cases  the  decomposition  goes  farther,  and  the 
metal  is  left  as  oxide : 

2  Pb(N03)2 +  2  PbO  +  4  N02  +  O2 

The  nitrates  are  especially  used  in  the  manufacture  of 
gunpowder,  sulfuric  acid,  nitric  acid,  and  as  a  fertilizer. 

Nitrous  acid  (HN02).  It  is  an  easy  matter  to  obtain  sodium 
nitrite  (NaN02)  by  heating  sodium  nitrate,  as  explained  in 
the  previous  paragraph.  Now  when  sodium  nitrite  is  treated 
with  an  acid,  such  as  sulfuric  acid,  it  is  decomposed  and 
nitrous  acid  is  set  free : 

NaN02  +  H2S04 >•  NaHS04  +  HN02 

The  acid  is  very  unstable,  however,  and  decomposes  into  water 
and  oxides  of  nitrogen.  Sodium  nitrite  is  used  in  the  manu- 
facture of  artificial  dyes. 

OXIDES  OF  NITROGEN 

The  most  important  of  the  oxides  of  nitrogen  are  the 
following : 

Nitrous  oxide  (N"2O) a  colorless  gas 

Nitric  oxide  (NO) a  colorless  gas 

Nitrogen  dioxide  (NO2)  ....  a  reddish-brown  gas 

Nitrogen  trioxide  (N2O3)     .     .     .  known  only  at  low  temperatures 

Nitrogen  pentoxide  (N2O5)       .     .  a  white  solid 


132 


FIRST  COUKSE 


CHEMISTRY 


Nitrous  oxide  (laughing   gas)  (N20).    This  gas  is  most 
readily  prepared  by  heating  ammonium  nitrate : 


NH4N03 


2  HO  +  NO 


It  is  colorless,  is  somewhat  soluble  in  water,  and  in  solution 
has  a  slightly  sweetish  taste.  When  inhaled,  it  produces  a 
kind  of  hysteria  (hence  the  name  laughing  gas)  and,  if 
taken  in  large  amounts,  insensibility  to  pain  and  uncon- 
sciousness. It  was  the  first  substance  to  be  used  as  an  anaes- 
thetic in  surgery,  and  is 
still  used  in  minor  opera- 
tions, such  as  those  of 
dentistry. 

Nitrous  oxide  is  a  very 
energetic  oxidizing  agent. 
Substances  such  as  car- 
bon, sulfur,  iron,  and 
phosphorus  burn  in  it 
almost  as  brilliantly  as  in 
oxygen,  forming  oxides 
and  setting  free  nitrogen. 
Evidently  the  oxygen  in  nitrous  oxide  is  not  held  in  very 
firm  combination  by  the  nitrogen. 

Nitric  oxide  (NO).  Nitric  oxide  is  most  conveniently 
prepared  by  the  action  of  nitric  acid  upon  copper : 

3  Cu  +  8  HN03 >•  3  Cu(NO8)a  +  2  NO  +  4  HaO 

The  metal  is  placed  in  the  flask  A  (Fig.  61),  and  the  acid 
added  slowly  through  the  funnel  tube  B.  The  gas  escapes 
through  C  and  is  collected  over  water.  Nitric  oxide  is  a  color- 
less gas  slightly  heavier  than  air.  Unlike  nitrous  oxide,  it 
does  not  part  with  its  oxygen  easily,  and  burning  substances 
introduced  into  this  gas  are  usually  extinguished. 


FIG.  61.  The  preparation  of  nitric  oxide 


COMPOUNDS  OF  NITROGEN 


133 


When  nitric  oxide  comes  into  contact  with  oxygen  or  air, 
it  at  once  combines  with  the  oxygen,  even  at  ordinary  tem- 
peratures, forming  a  reddish-brown  gas,  NO2,  which  is  called 
nitrogen  dioxide:  2  NQ  +  Q  _ 


2ND 


To  show  the  formation  of  nitrogen  dioxide  from  nitric  oxide 
and  oxygen,  a  tube  is  filled  with  the  oxide,  inverted  in  water, 
and  pure  oxygen  is  passed  into 
it,  as  shown  in  Fig.  62.  As  each 
bubble  of  oxygen  enters,  it  unites 
with  the  nitric  oxide  to  form  the 
reddish-brown  dioxide.  In  a  few 
minutes  the  color  fades  (because 
of  the  action  of  water  upon  the 
dioxide),  and  the  water  slowly 
rises  in  the  tube. 

Nitrogen  dioxide  (N02).  This 
gas,  as  we  have  just  seen,  is 
formed  by  allowing  nitric  oxide 
to  come  into  contact  with  oxy- 
gen. It  can  also  be  made  by 
heating  certain  nitrates,  such 

as  lead  nitrate  (p.  131).  It  is  a  reddish-brown  gas  of  un- 
pleasant odor,  and  is  poisonous  when  inhaled.  It  gives  up 
a  part  of  its  oxygen  to  burning  substances,  acting  as  an 
oxidizing  agent: 


FIG.  62.    The  formation  of  ni- 
trogen dioxide  from  nitric  oxide 
and  oxygen 


Nitrogen  tetroxide.  At  lower  temperatures  nitrogen  dioxide 
becomes  paler  in  color  and  condenses  to  a  pale-yellow  liquid.  It 
has  been  shown  that  this  paler  gas  has  the  formula  N204,  and  it 
is  called  nitrogen  tetroxide.  At  ordinary  temperatures  the  gas 
is  a  mixture  of  the  two,  and  we  may  express  this  relation  thus  : 


Nitrogen  dioxide,  2  NO, 

high  temperatures 


nitrogen  tetroxide,  N204 

low  temperatures 


134  FIKST  COUKSE  IN  CHEMISTRY 

Acid  anhydrides.  The  oxides  N2Og  (nitrogen  trioxide) 
and  N2O5  (nitrogen  pentoxide)  are  rarely  prepared  and 
need  not  be  separately  described.  They  bear  a  very  inter- 
esting relation  to  the  acids  of  nitrogen.  When  dissolved  in 
water  they  combine  with  the  water,  forming  acids  : 


NO  +  HO  -  >•  2  HN(X 

m       9  m  o 

Many  other  oxides  act  in  the  same  way,  combining  with 
water  to  form  an  acid.  Such  oxides  are  called  acid 
anhydrides. 

EXERCISES 

1.  Perfectly  dry  ammonia  does  not  affect  litmus  paper.   Explain. 

2.  Can  ammonia  be  dried  by  passing  the  gas  through  concen- 
trated sulf  uric  acid  ?    Explain. 

3.  Why  is  brine  used  in  the  manufacture  of  artificial  ice  ? 

4.  Discuss  the  energy  changes  which  take  place  in  the  manu- 
facture of  artificial  ice. 

5.  Write  the  equations  for  the  reactions  taking  place  when  am- 
monium hydroxide  is  neutralized  by  hydrochloric,  sulfuric,  and  nitric 
acids  respectively. 

6.  It  is  said  that  nitric  acid  is  formed  during  thunderstorms. 
How  would  you  account  for  its  formation  ? 

7.  What  does  the  word  ammonia  mean  (consult  dictionary)  ? 

8.  Why  is  nitric  acid  said  to  be  a  strong  acid  ? 

9.  What  are  the  properties  of  ammonia  that  make  it  suitable 
for  use  in  the  preparation  of  artificial  ice  ? 

10.  How  many  liters  of  ammonia  at  0°  and  760  mm.  pressure  will 
1  liter  of  water  dissolve?  Ans.  1298  liters.  What  would  this  volume 
of  ammonia  weigh?  Ans.  1000.5  g.  What  weight  of  ammonium 
chloride  would  be  necessary  to  prepare  it?  Ans.  3142.40  g. 

TOPICS  FOR  THEMES 

The  manufacture  of  ice.   (If  possible,  visit  an  ice  plant.) 
The  manufacture  of  nitric  acid  from  the  air  (Duncan,  Chemistry 
of  Commerce,  chapter  on  Fixation  of  Nitrogen). 
The  use  of  nitrous  oxide  as  an  anaesthetic. 


CHAPTER  XVI 
EQUILIBRIUM;    MASS  ACTION 

Reversible  reactions.  We  have  met  with  a  number  of 
reactions  which  are  especially  interesting  because  they  can 
go  in  either  direction.  Thus,  when  we  heat  mercuric  oxide, 
we  obtain  mercury  and  oxygen  ;  while  if  we  heat  mercury 
in  contact  with  oxygen,  we  obtain  mercuric  oxide  (p.  6). 
These  facts  are  represented  in  the  following  way: 

2  HgO  —  >•  2  Hg  +  O2 

In  a  similar  way  we  have  found  that  when  an  electric 
discharge  is  passed  through  a  mixture  of  nitrogen  and  hy- 
drogen, we  get  a  small  quantity  of  ammonia  (p.  123)  ;  yet 
when  the  discharge  is  passed  through  ammonia,  we  get 
hydrogen  and  nitrogen: 


Such  reactions  are  known  as  reversible  reactions. 

Equilibrium.  If  we  remember  that  the  materials  taking 
part  in  a  reaction  are  made  up  of  great  numbers  of  mole- 
cules all  of  which  are  in  rapid  motion  and  are  constantly 
changing  their  relations  to  each  other,  it  is  not  difficult  to 
see  why  some  molecules  should  be  decomposing  while  others 
are  forming.  In  time,  however,  a  condition  will  be  reached 
in  which  the  changes  in  the  one  direction  will  just  offset 
those  in  the  other.  The  average  percentage  of  each  material 
present  will  then  remain  unchanged.  This  condition  of 

135 


136  FIEST  COUESE  IN  CHEMISTRY 

affairs  is  called  equilibrium.  Thus,  ammonia,  hydrogen,  and 
nitrogen  come  to  equilibrium  in  the  presence  of  electric 
discharge  when  there  is  about  7  per  cent  of  ammonia  and 
93  per  cent  of  uncombined  nitrogen  and  hydrogen. 

Mass  action.  Suppose,  when  equilibrium  has  been 
reached,  we  add  an  additional  quantity  of  one  of  the  act- 
ing substances  —  say  hydrogen  in  the  case  just  mentioned. 
This  will  make  it  easier  for  the  nitrogen  to  act  upon  the 
hydrogen,  for  the  two  kinds  of  molecules  will  now  meet 
more  frequently.  It  will  not  at  all  affect  the  rate  at  which 
ammonia  is  decomposing.  The  net  effect  will  therefore  be 
to  bring  about  a  new  equilibrium  in  which  a  larger  per- 
centage of  ammonia  is  present.  The  effect  produced  by  an 
excess  of  one  of  the  reacting  materials  is  called  mass  action. 

Changing  an  equilibrium  to  a  completed  reaction.  If  we 
were  to  withdraw  the  ammonia  as  fast  as  it  is  formed,  be- 
fore it  has  time  to  decompose,  the  reaction  ought  to  go  on 
until  either  the  hydrogen  or  the  nitrogen  is  used  up.  This 
is  just  what  takes  place  if  during  the  discharge  the  gases 
are  inclosed  over  water  containing  acid,  with  which  the 
ammonia  combines. 

The  point  of  equilibrium  can  therefore  be  changed  or  the 
equilibrium  converted  into  a  completed  reaction  by  changing 
the  acting  mass  of  the  substances  taking  part  in  the  reaction. 

Equilibrium  in  solution.  In  aqueous  solution  the  mole- 
cules of  an  electrolyte  keep  dissociating  into  ions,  while  the 
ions,  on  meeting,  recombine  to  form  molecules,  the  result 
being  an  equilibrium  between  the  two  conditions,  thus  : 


If  we  mix  two  electrolytes,  the  equilibrium  reached  be- 
comes much  more  complicated,  for  any  positive  ion  may 
unite  with  any  negative  one.  At  equilibrium  all  possible 


EQUILIBRIUM;  MASS  ACTION 


13T 


ions  and  combinations  of  ions  will  be  present.  Thus,  when 
we  mix  sodium  nitrate  and  sulfuric  acid  in  the  prepara- 
tion of  nitric  acid,  we  have  present  the  ions  Na+,  NO3~,  H+, 
and  SO4~~,  together  with  the  molecules  NaNO3,  Na0SO4, 
NaHSO4,  HNO3,  and  H2SO4. 

Completion  of  reactions  in  solution.  The  chemist  makes 
use  of  reactions  to  secure  various  compounds  in  pure 
condition,  and  he  wishes  the 
yield  to  be  as  large  as  possible. 
Reactions  which  stop  short 
of  completion  and  end  in  an 
equilibrium  are  not  suited  to 
manufacturing  purposes,  un- 
less means  can  be  found  to 
break  up  the  condition  of 
equilibrium  and  bring  the 
reaction  to  a  definite  conclu- 
sion. There  are  three  condi- 
tions under  which  this  may 
be  accomplished. 

1.  A    volatile  gas   may    be 
formed.     If    the    reaction    is 
conducted    under   such   con- 
ditions that  one  of  the  prod- 
ucts is  a  gas  insoluble  in  the  solvent,  the  gas  will  make  its 
escape  as  fast  as  it  is  formed,  and  this  action  will  con- 
tinue until  one  or  the  other  of  the  ions  taking  part  in  its 
formation  is  used  up. 

Thus,  when  we  mix  sulfuric  acid  and  sodium  nitrate,  no 
visible  reaction  takes  place..  But  if  we  heat  the  mixture 
above  the  boiling  point  of  nitric  acid,  all  of  this  substance 
formed  in  the  equilibrium  between  the  H+  and  the  NO3~ 
ions  is  converted  into  a  gas  insoluble  in  sulfuric  acid.  The 


FIG.  03.    Precipitation  of  silver 
chloride 


138  FIRST  COURSE  IN  CHEMISTRY 

nitric  acid  distills  away  until  the  NOg~  ions  are  used  up.  We 
then  have  a  completed  reaction  expressed  in  the  equation 

NaN03  4-  H2S04  — >•  NaHSO4  +  HNO3 

It  is  in  this  way  that  most  acids  are  prepared.  Their  salts 
are  heated  with  some  acid  of  high  boiling  point,  usually 
with  sulfuric  acid,  which  boils  at  338°. 

2.  An  insoluble  solid  may  be  formed.    When  hydrochloric 
acid  (HC1)  and  silver  nitrate  (AgNOa)  are  brought  to- 
gether   in    solution,    we    have   in    addition   to   these   two 
compounds  the  ions  H+,  Cl~,  Ag+,  NO8~,  and  the  new 
combinations  HNO3  and  AgCl.  One  of  these,  namely,  silver 
chloride  (AgCl),  is  insoluble  in  water  and  as  fast  as  formed 
separates   from  the  solution  as   a  curdy  white  precipitate 
(Fig.  63).    The  reaction  therefore  continues  till  either  the 
Ag+  or  the  NO3~  is  used  up,  the  completed  equation  being 

H+  H-  Cl-  +  Ag+  4-  N03~  •  ->•  H+  +  NO3~  +  AgCl 

3.  Two  different  ions  may  unite  to  form  an  undissociated 
molecule.    When  we  bring  together  sodium  hydroxide  and 
hydrochloric  acid  in  solution,  we  have  the  ions  H+,  Cl~, 
Na+,  and  OH~.  The  H+  ions  and  the  OH~  ions  unite  to  form 
molecules  of  water  which  do  not  again  part  into  ions  save 
to  a  very  slight  extent.    This  leaves  only  the  ions  of  NaCl 
in  solution,  the  equation  being 

Na+  +  OH-  +  H+  +  Cl-  •  -»-  H2O  +  Na+  4  Cl~ 

Neutralization  is  practically  a  completed  reaction  because 
water  is  so  little  ionized. 

Hydrolysis.  While  water  is  very  little  ionized,  neverthe- 
less it  forms  some  ions.  Moreover,  when  a  salt  is  dissolved 
in  water  to  form  a  dilute  solution,  the  relative  mass  of 
the  water  is  very  great.  The  reaction  of  neutralization  is 


EQUILIBRIUM;  MASS  ACTION  139 

therefore  reversed  to  a  slight  extent,  forming  a  small  amount 
of  free  base  and  of  free  acid,  thus : 

NaN02  +  H2O  -<->  NaOH  +  HNO2 

A  reaction  of  this  kind,  in  which  water  acts  upon  a  salt  to 
form  a  base  and  an  acid,  is  called  hydrolysis.  If  the  base 
formed  in  hydrolysis  is  very  weak  and  the  acid  is  strong, 
the  solution  will  turn  blue  litmus  red,  as  is  true  with  all 
salts  of  aluminium.  If  the  base  is  very  strong  and  the  acid 
weak,  the  solution  will  turn  red  litmus  blue,  as  is  the  case 
with  many  salts  of  sodium.  If  both  the  acid  and  the  base 
are  weak,  then  the  compound  may  be  completely  hydrolyzed. 

EXERCISES 

1.  Can  you  mention  any  reversible  reactions,  other  than  those 
given  in  this  chapter? 

2.  Suggest  a  method  for  the  preparation  of  hydrogen  chloride. 

3.  Would  silver  nitrate  produce  a  precipitate  when  added  to  a 
solution  of  sodium  chloride  (common  salt)  ?    If  so,  how  would  the 
precipitate  compare  in  composition  with  that  produced  when  silver 
nitrate  is  added  to  hydrochloric  acid  ? 

4.  Barium  sulfate  (BaSO4)  is  a  white  insoluble  compound  much 
used  as  a  pigment  in  making  paints.  Suggest  a  method  for  preparing  it. 

5.  Is  the  reaction  NH3  +  H2O >-NII4OH  reversible?   If  so, 

state  the  conditions  under  which  it  will  go  in  each  direction. 

6.  Is  the  reaction  expressed  by  the  equation  2  H2  +  O2 >-  2  H2O 

reversible  ?   If  so,  state  the  conditions  under  which  it  will  go  in  each 
direction. 

7.  Carbonic  acid  is  a  very  weak  acid,  while  sodium  hydroxide  is 
a  strong  base.    How  will  a  solution  of  sodium  carbonate  act  towards 
litmus  paper? 


CHAPTER  XVII 
SULFUR  AND  ITS  COMPOUNDS 

Occurrence.  The  element  sulfur  has  been  known  from 
the  earliest  times.  It  is  widely  distributed  in  nature,  and 
occurs  in  large  quantities  in  the  uncombined  form,  espe- 
cially in  the  neighborhood  of  volcanoes.  Sicily  has  long 
been  famous  for  its  sulfur  mines,  and  in  more  recent  years 


FIG.  64.   The  flow  of  liquid  sulfur  from  a  well  (pipe  A)  in  Louisiana 

large  deposits  have  been  found  in  Louisiana.  It  is  occa- 
sionally found  in  well-formed  crystals. 

In  combination,  sulfur  occurs  abundantly  in  the  form 
of  sulfides  and  sulfates.  In  smaller  amounts  it  is  found  in  a 
great  variety  of  minerals  and  is  a  constituent  of  many  vege- 
table and  animal  substances,  especially  of  the  yolk  of  eggs. 

Extraction  of  sulfur.  In  Louisiana  the  sulfur  occurs  in 
deposits  far  underground  and  covered  with  quicksand  so 
that  it  cannot  be  mined.  One  of  these  deposits  lies  at  a 

140 


SULFUR  AND  ITS  COMPOUNDS 


141 


depth  of  700  feet,  is  circular  in  shape,  and  is  about  half  a 
mile  in  diameter  and  500  feet  in  thickness.  Wells  are 
drilled  into  the  deposit,  and  superheated  water  (above  160°) 
is  forced  down  through  suitable  pipes.  The  hot  water  melts 


FIG.  65.    Loading  a  block  of  sulfur  on  a  car  in  Louisiana 

the  sulfur,  which  is  then  forced  up  a  separate  pipe  by  com- 
pressed air  (Fig.  64).  The  liquid  sulfur  then  solidifies  in 
very  large  blocks  (Fig.  65).  A  single  well  has  produced 
500  tons  daily,  and  the  product  is  99.5  per  cent  pure.  About 
250,000  tons  are  produced  annually  from  this  deposit. 

In  Sicily  a  very  simple  but  wasteful  method  is  used  to 
separate  sulfur  from  the  rock 
and    earthy   materials    with 
which  it  is  mixed.    The  ma- 
terial is  piled  up  in   heaps 
and  set  on  fire,  and  the  heat 
from  the  burning  of  a  part 
of    the    sulfur    serves    to 
melt  another  portion,  which 
collects  as  a  liquid  at  the 

bottom  of  the  pile.  This  is  drained  off  and  purified  by 
distillation  in  a  retort-shaped  vessel  A  (Fig.  66),  the  exit 
tube  of  which  opens  into  a  cooling  chamber,  B,  of  brickwork. 


FIG.  66.   A  sulfur  still 


142 


FIEST  COURSE  IN  CHEMISTRY 


When  the  sulfur  vapor  first  enters  the  cold  chamber  it 
condenses  as  a  fine  crystalline  powder  called  flowers  of  sulfur. 
As  the  condensing  chamber  becomes  warm,  the  sulfur  con- 
denses as  a  liquid  and  is  drawn  off  into  cylindrical  molds, 
the  product  being  called  roll  sulfur  or  lriinxf'>ne. 

Varieties  of  sulfur.  Sulfur  exists  in  a  number  of  quite 
different  forms.  Several  other  elements  occur  in  a  number 
of  different  forms,  but  the  forms  of  sulfur  are  unusually 
numerous  and  are  easy  to  obtain.  The  best-known  forms 

are  the  following: 

1.  Ordinary,  or  rhombic, 
sulfur.  When  sulfur  crys- 
tallizes from  solution  in 
liquids  (notably  from  car- 
bon disulfide)  it  is  obtained 
in  compact  yellow  crystals 
which  melt  at  114.5°  and 
have  a  density  of  2.06. 
This  is  called  rhombic  sul- 
fur (Fig.  67),  and  roll 
sulfur  is  composed  largely 
of  this  variety. 

2.  Prismatic,  or  monoclinic,  sulfur.  When  melted  sulfur  is 
allowed  to  cool  until  a  part  of  the  liquid  has  solidified,  and 
the  remaining  liquid  is  then  poured  off,  it  is  found  that  the 
solid  sulfur  remaining  in  the  vessel  is  in  the  form  of  fine 
needle-shaped  crystals,  which  melt  at  119°  and  have  a  den- 
sity of  1.96.    The  needle-shaped  form  is  called   nmnndinic 
sulfur.    At  all  temperatures  below  96°  the  needle-shaped 
crystals  break  up  more  or  less  rapidly  into  little  crystals 
of  the  rhombic  variety. 

3.  Plastic  sulfur.   When  boiling  sulfur  is  poured  into  cold 
water  it  assumes  a  gummy,  doughlike  form,  which  is  quite 


FIG.  67.    Natural  crystals  of  rhombic 
sulfur 


SULFUR  AND  ITS  COMPOUNDS 


143 


elastic.  This  can  be  seen  in  a  very  striking  manner  by  dis- 
tilling sulfur  from  a  small,  short-necked  retort  (such  as  is 
represented  in  Fig.  68)  and  allowing  the  liquid  to  run 
directly  into  water.  In  a  few  clays  it  becomes  brittle  and 
in  part  passes  over  into  ordinary  rhombic  sulfur. 

4.  White,  or  amorphous,  sulfur.  If  freshly  prepared  plastic 
sulfur  is  treated  with  carbon  disulfide,  the  rhombic  sulfur 
in  it  dissolves  and  a  nearly  colorless  residue  remains  which 
is  not  crystalline.  This  is  called  white,  or  amorphous,  sulfur. 
Ordinary  flowers  of  sulfur  con- 
sist of  a  mixture  of  rhombic  crys- 
tals and  amorphous  particles. 

Properties  of  ordinary  sulfur. 
Sulfur  is  a  pale-yellow  crystal- 
line solid,  without  marked  taste 
and  with  but  a  faint  odor.  It 
is  insoluble  in  water,  but  is 
freely  soluble  in  a  few  liquids, 
notably  in  carbon  disulfide.  It 
melts  at  114.5°.  Just  above  the 
melting  point  it  forms  a  rather 
thin,  straw-colored  liquid.  As 
the  temperature  is  raised,  this  liquid  turns  darker  in  color 
and  becomes  thicker,  until  at  about  235°  it  is  almost  black 
and  is  so  thick  that  the  vessel  containing  it  can  be  inverted 
without  danger  of  the  liquid  running  out.  At  higher  tem- 
peratures it  becomes  thin  once  more,  and  boils  at  444.6°, 
forming  a  yellowish  vapor.  When  the  vapor  cools,  the  same 
changes  take  place  in  reverse  order. 

Chemical  conduct  of  sulfur.  Sulfur  burns  in  oxygen  or  in 
the  air  with  a  pale-blue  flame,  forming  sulfur  dioxide  (SO2). 
Most  metals  when  heated  with  sulfur  combine  directly  with 
it,  forming  metallic  sulfides.  In  some  cases  the  action  is  so 


FIG.  68.  The  preparation  of 
plastic  sulfur 


144 


FIRST  COURSE  IN  CHEMISTRY 


energetic  that  the  mass  becomes  incandescent,  as  is  the  case 
with  iron.  This  conduct  recalls  the  action  of  oxygen  upon 
metals,  and  in  general  the  metals  which  combine  readily 
with  oxygen  are  apt  to  combine  with  sulfur. 

Uses  of  sulfur.  Large  quantities  of  sulfur  are  used  in 
the  manufacture  of  gunpowder,  vulcanized  rubber,  carbon 
disulfide,  sulfur  dioxide,  sulfuric  acid,  and  salts  of  various 


FIG.  69.    Spraying  an  orchard  of  fruit  trees  with  lime-sulfur  spray 

kinds.    It  is  also  used  extensively  in  the  manufacture  of 
insecticides  for  use  in  orchards  and  vineyards. 

Lime-sulfur  spray.  The  chief  sulfur  insecticide  is  known 
as  lime-sulfur  spray.  It  is  made  by  boiling  sulfur  with 
slaked  lime,  by  which  process  a  deep-red  liquor  is  obtained, 
consisting  essentially  of  a  solution  of  sulndes  of  calcium 
(CaS4  and  CaS5).  The  liquid  is  a  very  efficient  insecticide, 
particularly  for  scale,  and  it  is  also  a  fungicide.  Large 
quantities  of  it  are  used  for  spraying  fruit  trees  (Fig.  69). 


SULFUR  AND  ITS  COMPOUNDS  145 

Vulcanized  rubber.  Natural  rubber  (caoutchouc)  is  very 
soft  and  elastic,  and  becomes  fluid  at  about  120°.  To  ren- 
der it  suitable  for  most  purposes  it  must  be  vulcanized. 
This  process  consists  in  working  into  the  caoutchouc  about 
2  to  3  per  cent  of  sulfur,  though  for  some  purposes  other 
materials,  such  as  zinc  oxide  and  antimony  sulfide,  are  used. 
The  sulfur  is  worked  into  the  warm  rubber  or  the  rubber 
is  immersed  in  liquids  containing  sulfur.  Vulcanized  rub- 
ber is  more  elastic  than  caoutchouc  and  much  less  sticky. 

When  a  larger  percentage  of  sulfur  is  added  and  the 
product  is  heated  somewhat  higher,  a  black,  horny  material 
is  obtained  called  hard  rubber,  vulcanite,  or  ebonite.  It  is 
used  for  making  such  articles  as  combs,  buttons,  fountain 
pens,  and  electrical  insulators. 

The  crude  rubber  is  obtained  from  the  milky  sap  of  certain 
species  of  trees  growing  in  tropical  countries,  particularly  in 
Brazil.  The  sap  is  collected  in  buckets,  and  paddle-shaped 
sticks  are  dipped  in  and  then  dried  over  a  fire.  The  process  is 
repeated  until  a  lump  of  crude  rubber  has  been  collected,  and 
this  is  then  loosened  from  the  stick  and  sent  to  the  market. 


COMPOUNDS  OF  SULFUR  WITH  HYDROGEN 

Hydrogen  sulfide  (H2S).  This  gas  is  found  in  the  vapors 
issuing  from  volcanoes,  and  in  solution  in  many  natural 
waters.  It  is  formed  when  organic  matter  containing  sulfur 
undergoes  decay,  just  as,  under  similar  circumstances, 
ammonia  is  formed  from  nitrogenous  matter. 

Preparation.  Since  hydrogen  sulfide  is  a  gas  which  is  a 
little  soluble  in  water,  it  can  be  prepared  by  treating  a  sul- 
fide with  an  acid  (p.  138).  Iron  sulfide  (FeS)  is  usually 
employed : 

FeS  +  2  HC1 >•  FeCl2  +  H2S 


146 


FIRST  COURSE  IN  CHEMISTRY 


A  convenient  apparatus  is  shown  in  Fig.  70.  A  few  lumps 
of  iron  sullide  are  placed  in  the  bottle  A  and  dilute  acid  is 
added  a  little  at  a  time  through  the  funnel  tube  B,  the  gas 
escaping  through  the  tube  C. 

Properties.  Hydrogen  sulfide  is  a  colorless  gas  having  a 
weak,  disagreeable  taste  and  a  most  offensive  odor  suggest- 
ing rotten  eggs.  At  ordinary  temperatures  it  is  but  spar- 
ingly soluble  in  water;  in  boiling  water  it  is  not  soluble  at 

all.  When  inhaled  in  concentrated 
form  it  acts  as  a  violent  poison, 
and  even  when  much  diluted  with 
air,  produces  headache,  dizziness, 
and  nausea.  It  is  a  little  heavier 
than  air,  having  a  density  of  1.18. 
Chemical  conduct.  The  most  im- 
portant chemical  properties  of  hy- 
drogen sulfide  are  the  following : 
1.  Acid  properties.  When  dis- 
solved in  water,  hydrogen  sulfide 
acts  as  a  weak  acid,  the  solution 
being  sometimes  called  hydro 
sulfuric  acid.  The  solution  turns  blue  litmus  red,  and 
neutralizes  bases,  forming  salts  called  mlfides. 

2.  Action  with  oxygen.  The  elements  composing  hydrogen 
sulfide  have  each  a  strong  tendency  to  combine  with  oxy- 
gen, and  are  not  held  together  very  firmly.  Consequently 
the  gas  burns  readily  in  oxygen  or  air  according  to  the 
equation  .  2  H2S  +  3  O2 +  2  H2O  +  2  SO2 

When  there  is  not  enough  oxygen  for  both  the  sulfur  and 
the  hydrogen,  the  latter  element  combines  with  the  oxygen, 
and  the  sulfur  is  set  free : 


FIG.  70.   The  preparation  of 
hydrogen  sulfide 


SULFUR  AND  ITS  COMPOUNDS 


147 


3.  Reducing  action.    Owing  to  the  ease  with  which  hydro- 
gen sulfide  decomposes  and  the  strong  tendency  of  both 
sulfur  and  hydrogen  to  combine   with   oxygen,   the   sub- 
stance is  a  strong  reducing  agent. 

4.  Action  on  metals.    Hydrogen  sulfide  acts  upon  many 
metals,  forming  sulfides.    Silver  sulfide  (Ag2S)  is  black,  and 
it  is  owing  to  traces  of 

hydrogen  sulfide  in  the 
air  that  silver  objects 
tarnish. 

Sulfur  springs.  The 
waters  of  many  natural 
springs  hold  hydrogen 
sulfide  in  solution,  as  is 
indicated  by  their  strong 
odor  and  the  way  in 
which  they  will  blacken 
a  silver  coin.  When  the 
water  reaches  the  air, 
the  hydrogen  sulfide  is 
slowly  oxidized,  with  the  liberation  of  sulfur,  which  often 
deposits  about  the  borders  of  the  spring. 

Salts  of  hydrosulfuric  acid  ;  sulfides.  The  salts  of  hydro- 
sulfuric  acid  (called  sulfides)  form  an  important  class  of 
salts.  Many  of  them  are  found  abundantly  in  nature,  and 
some  of  them  are  important  ores. 

Uses  of  the  sulfides  in  analysis.  Most  of  the  sulfides  are 
insoluble  in  water,  and  some  of  them  are  insoluble  in  acids. 
Consequently,  when  hydrogen  sulfide  is  passed  into  a  solu- 
tion of  a  salt,  it  often  happens  that  a  sulfide  is  precipitated. 
With  copper  chloride  the  equation  is 


FIG.  71. 


The  preparation  of  insoluble 
sulfides 


CuCl2  +  H2S  — >-  CuS  +  2  HC1 


148  FIRST  COURSE  IN  CHEMISTRY 

Because  of  the  fact  that  some  metals  are  precipitated  in 
this  way  as  sulfides  while  others  are  not,  hydrogen  sulfide 
is  extensively  used  in  the  separation  of  the  metals  in  the 
laboratory. 

The  sulfides  are  prepared  in  the  laboratory  by  passing  hy- 
drogen sulfide  through  solutions  of  the  salts  of  the  metals  as 
shown  in  Fig.  71.  The  hydrogen  sulfide  generated  in  flask  A , 
is  passed  through  bottles  B  and  C,  containing,  say,  solutions 
of  silver  nitrate  and  arsenic  chloride  respectively.  As  the 
gas  bubbles  through  the  solutions  there  is  formed  black  silver 
sulfide  in  B  and  yellow  arsenic  sulfide  in  C. 

OXIDES  OF  SULFUR 

Sulfur  forms  two  well-known  compounds  with  oxygen : 
sulfur  dioxide  (SO2),  sometimes  called  sulfurous  anhy- 
dride; and  sulfur  trioxide  (SO8),  frequently  called  sulfuric 
anhydride. 

Sulfur  dioxide  (S02).  Sulfur  dioxide  often  occurs  in 
nature  in  the  gases  issuing  from  volcanoes,  and  in  solution 
in  the  water  of  many  springs.  It  is  likely  to  be  found 
wherever  sulfur  compounds  are  undergoing  oxidation. 

Preparation.  Two  general  ways  may  be  mentioned  for 
the  preparation  of  sulfur  dioxide : 

1.  By  the  combustion  of  sulfur.    Sulfur  dioxide  is  readily 
formed  when  sulfur  or  certain  compounds  containing  sulfur, 
such  as  the  metal  sulfides,  are  heated  in  air  or  oxygen : 

2  ZnS  +  3  O2  — *  2  ZnO  +  2  SO2 

2.  By  the  reduction  of  sulfuric  acid.    When  concentrated 
sulfuric  acid  is  heated  with  certain  metals,  such  as  copper, 
part  of  the  acid  is  changed  into  copper  sulfate  and  part  is 


SULFUR  AND  ITS  COMPOUNDS 


149 


reduced  to  sulfurous  acid.    The  latter  then  decomposes  into 
sulfur  dioxide  and  water,  the  complete  equation  being 


2  HSO 


CuSO  +  SO  +  2  HO 


Properties.  Sulfur  dioxide  is  a  colorless  gas,  which  at 
ordinary  temperatures  is  2.2  times  as  heavy  as  air.  It  has 
a  peculiar,  irritating  odor.  The  gas  is  very  soluble  in 
water,  1  volume  of  water  dissolving  80  volumes  of  the  gas 
under  standard  conditions.  It  is  easily 
condensed  to  a  colorless  liquid,  and  can 
be  purchased  in  this  condition,  stored 
in  strong  bottles  or  in  metal  cylinders 
(Fig.  72). 

Chemical  conduct.  Sulfur  dioxide  has 
a  marked  tendency  to  combine  with  other 
substances,  and  is  therefore  an  active  sub- 
stance chemically.  It  combines  with  oxy- 
gen gas,  but  not  very  easily.  It  can, 
however,  take  oxygen  away  from  some 
other  substances,  and  is  therefore  a  good 

reducing  agent.    Its  most  marked  chem- 

.      ..         ,  ...A  ,  .          FIG.  72.   A  cylinder 

ical  property  is  its    ability  to  combine      Of  sulfur  dioxide 

with  water. 

Sulfurous  acid  (H2S03).  When  sulfur  dioxide  is  passed 
into  water,  it  combines  chemically  with  it  to  form  sulfu- 
rous acid  (H2SOg).  It  is  impossible  to  prepare  this  acid  in 
pure  form,  as  it  breaks  down  very  easily  into  water  and  sul- 
fur dioxide.  The  reaction  is  therefore  reversible,  and  is 
expressed  by  the  equation 


HO  +  SO, 


HSO, 


Solutions  of  the  acid  in  water  are  often  prepared  and  have 
a  number  of  interesting  properties  and  commercial  uses. 


150 


FIRST  COURSE  IN  CHEMISTRY 


1.  Acid  properties.    The  solution  has  all  the  properties 
typical  of  a  very  weak  acid.    When  neutralized  by  bases, 
sulfurous  acid  yields  a  series  of  salts  called  sulfites,  most 
of  which  are  insoluble  in  water. 

2.  Reducing  properties.    Solutions  of  sulfurous  acid  act  as 
good  reducing  agents.    This  is  due  to  the  fact  that  sulfurous 
acid  has  the  power  of  taking  up  oxygen  from  the  air  or 
from  substances  rich  in  oxygen,  and  is  changed  by  this 

reaction  into  sulfuric  acid : 

2HJS(X  +  (X — *-2HSO. 


3.  Bleaching  properties.  Sulfurous  acid 
has  strong  bleaching  properties,  acting 
upon  many  colored  substances  in  such 
a  way  as  to  destroy  their  color.  It  is 
on  this  account  used  to  bleach  paper, 
straw  goods,  and  even  such  foods  as 
canned  corn  and  dried  fruits.  As  a  rule 
the  bleaching  is  not  permanent. 


FIG.  73.    Bleaching 

a  flower  with  sulfur 

dioxide 


The  bleaching  properties  of  sulfurous  acid  may  be  shown 
by  bringing  a  small  dish  of  burning  sulfur  under  a  bell  jar 
(Fig.  73)  in  which  has  been  placed  some  highly  colored  flower 
moistened  with  water.  Straw  hats  may  be  cleaned  and  bright- 
ened in  a  similar  way. 

4.  Antiseptic  properties.  Sulfurous  acid  has  marked  anti- 
septic properties,  and  on  this  account  has  the  power  of  ar- 
resting fermentation.  It  is  therefore  used  in  certain  foods 
containing  sugars,  such  as  sweet  cider,  canned  corn,  and 
dried  fruits,  serving  not  only  as  a  preservative  but  also  as  a 
bleaching  agent  to  remove  objectionable  colors.  Whether  or 
not  its  use  in  foods  should  be  permitted,  is  a  much  debated 
question. 


SULFUR  AND  ITS  COMPOUNDS 


151 


Salts  of  sulfurous  acid ;  sulfites.  The  sulfites  are  solid 
compounds  and,  like  sulfurous  acid,  have  the  power  of  tak- 
'ing  up  oxygen  very  readily,  and  are  good  reducing  agents. 
On  account  of  this  tendency,  commercial  sulfites  are  often 
contaminated  with  sulphates. 

Sulfur  trioxide  (S03).  When  sulfur  dioxide  and  oxygen 
are  heated  together  at  a  rather  high  temperature,  a  small 
amount  of  sulfur  trioxide  (SO3)  is  formed,  but  the  reaction 

A 


FIG.  74.   The  preparation  of  sulfur  trioxide 

is  slow  and  incomplete.  If,  however,  the  heating  takes 
place  in  the  presence  of  very  fine  platinum  dust,  the  reaction 
is  rapid  and  nearly  complete  : 


Experimental  preparation  of  sulfur  trioxide.  The  experiment 
can  be  performed  by  the  use  o£  the  apparatus  shown  in  Fig.  74, 
the  fine  platinum  being  secured  by  moistening  asbestos  fiber 
with  a  solution  of  chloroplatinic  acid  and  igniting  it  in  a  flame. 
The  fiber,  covered  with  fine  platinum,  is  placed  in  a  tube  of 
hard  glass  A,  which  is  then  heated  with  a  burner  to  about  400°, 
while  sulfur  dioxide  and  air  are  passed  into  the  tube  through 
the  drying  bottles  B  and  C.  Union  takes  place  at  once,  and 
the  strongly  fuming  sulfur  trioxide  escapes  from  the  jet  at  the 
end  of  the  tube,  or  it  may  be  condensed  by  surrounding  the 
receiving  tube  D  with  a  freezing  mixture. 


152  FIRST  COURSE  IN  CHEMISTRY 

Properties  of  sulfur  trioxide.  Sulfur  trioxide  is  a  color- 
less liquid  which  solidifies  at  about  15°  and  boils  at  46°. 
A  trace  of  moisture  causes  it  to  solidify  into  a  mass  of  silky 
white  crystals  somewhat  resembling  asbestos  fiber  in  ap- 
pearance. In  contact  with  the  air  it  fumes  strongly,  and 
when  thrown  upon  water  it  dissolves  with  a  hissing  sound 
and  the  liberation  of  a  great  deal  of  heat.  The  product  of 
this  reaction  is  sulfuric  acid,  so  that  sulfur  trioxide  is  the 
anhydride  of  that  acid : 

S08  +  H20— >H2S04 

Catalysis.  It  has  been  found  that  many  chemical  reac- 
tions, such  as  the  union  of  sulfur  dioxide  with  oxygen,  are 
much  influenced  by  the  presence  of  substances  which  do 
not  themselves  seem  to  take  a  part  in  the  reaction  and 
are  left  apparently  unchanged  after  it  has  ceased.  These 
reactions  go  on  very  slowly  under  ordinary  circumstances, 
but  are  greatly  hastened  by  the  presence  of  the  foreign  sub- 
stance. Substances  which  increase  the  speed  of  reactions  in 
this  way  are  said  to  act  as  catalytic  agents  or  catalyzers,  and 
the  action  is  called  catalysis.  Just  how  the  action  is  brought 
about  is  not  well  understood,  but  the  part  played  by  the 
catalyzer  is  no  doubt  different  in  different  cases. 

Examples  of  catalysis.  Oxygen  and  hydrogen  combine  with 
each  other  at  ordinary  temperatures  in  the  presence  of  platinum 
powder,  while  if  no  catalytic  agent  is  present,  they  do  not  com- 
bine in  appreciable  quantities  until  a  much  higher  temperature 
is  reached.  Potassium  chlorate,  when  heated  with  manganese 
dioxide,  gives  up  its  oxygen  at  a  much  lower  temperature 
than  when  heated  alone  (p.  16).  Hydrogen  dioxide  decomposes 
very  rapidly  when  powdered  manganese  dioxide  is  sifted  into 
its  concentrated  solution.  In  the  Haber  process  for  preparing 
ammonia,  metallic  iron  promotes  the  reaction  between  nitrogen 
and  hydrogen. 


SULFUR  AND  ITS  COMPOUNDS  153 

On  the  other  hand,  the  catalytic  agent  sometimes  retards 
chemical  action  ;  for  example,  a  solution  of  hydrogen  dioxide 
decomposes  more  slowly  when  it  contains  a  little  phosphoric 
acid  or  acetanilide  than  when  perfectly  pure.  It  is  probable 
that  many  of  the  chemical  transformations  in  physiological 
processes,  such  as  digestion,  are  assisted  by  certain  substances 
acting  as  catalytic  agents.  The  principle  of  catalysis  is  there- 
fore very  important. 

Sulfuric  acid  (oil  of  vitriol)  (H2SOJ.  Sulfuric  acid  is  one 
of  the  most  important  of  all  manufactured  chemicals.  Not 
only  is  it  one  of  the  most  common  reagents  in  the  labora- 
tory, but  enormous  quantities  of  it  are  used  in  many  of 
the  industries,  especially  in  the  refining  of  petroleum  and 
the  manufacture  of  nitroglycerin,  hydrochloric  and  nitric 
acids,  sodium  carbonate,  and  phosphate  fertilizers. 

Manufacture  of  sulfuric  acid.  Sulfuric  acid  can  be  made 
at  low  cost,  and  is  the  cheapest  of  the  commercial  acids. 
Two  general  methods  are  used  in  its  manufacture. 

1.  Contact  process.     In  this  process    sulfur  trioxide    is 
made  from  sulfur  dioxide  and  oxygen,  as  explained  under 
Fig.  74.    The  two  gases  are  conducted  through  iron  tubes 
filled  with  some  porous  material,  such  as  asbestos  or  sodium 
sulfate,  through  which  is  interspersed  a  suitable  catalyzer, 
such  as  iron  oxide  or  platinum.    The  sulfur  trioxide  so 
formed  reacts  with  water  to  form  sulfuric  acid : 

S03+H20 — ^H2S04 

The  contact  process  is  only  used  when  the  concentrated 
acid  is  desired. 

2.  Chamber  process.    The  older  method  of  manufacture, 
exclusively  employed  until  recent  years  and  still  the  most 
important  process,  is  much  more  complicated.    The  conver- 
sion of  water,  sulfur  dioxide,  and  oxygen  into  sulfuric  acid 


154  FIKST  COURSE  IN  CHEMISTRY 

is  accomplished  by  the  catalytic  action  of  oxides  of  nitrogen. 
Since  these  oxides  are  gases,  it  is  difficult  to  prevent  their 
escape,  and  very  elaborate  precautions  have  to  be  taken  to 
reduce  the  loss  as  much  as  possible.  The  reactions  are 
brought  about  in  large,  lead-lined  chambers,  into  which 
oxides  of  nitrogen,  sulfur  dioxide,  steam,  and  air  are  intro- 
duced in  suitable  proportions. 

The  dilute  acid  resulting  collects  upon  the  floor  of  the 
lead  chambers.  It  is  drawn  off,  and  in  this  form  serves  for 
many  purposes,  such  as  the  manufacture  of  fertilizers.  The 
pure  concentrated  acid  can  be  prepared  from  the  dilute  acid, 
but  the  process  is  costly,  so  that  it  sometimes  is  cheaper 
to  prepare  this  form  of  acid  by  the  contact  process. 

Properties.  Pure  anhydrous  sulfuric  acid,  more  properly 
termed  hydrogen  sulfate,  is  a  colorless,  oily  liquid  nearly 
twice  as  heavy  as  water.  The  ordinary  concentrated  acid 
contains  about  2  per  cent  of  water,  has  a  density  of  1.84, 
and  boils  at  338°.  It  is  sometimes  called  oil  of  vitriol,  since 
it  was  formerly  made  by  distilling  a  mixture  of  substances, 
one  of  which  was  called  green  vitriol. 

Chemical  conduct.  Sulfuric  acid  possesses  chemical  prop- 
erties which  make  it  one  of  the  most  important  of  chemical 
substances. 

1.  Action  as  an  acid.   In  dilute  solution  sulfuric  acid  acts 
as  a  very  strong  acid,  turning  blue  litmus  red  and  forming 
salts  with  oxides  and  hydroxides. 

2.  Action  as  an  oxidizing  agent.    Sulfuric  acid  contains  a 
large  percentage  of  oxygen  and  is,  like  nitric  acid,  a  very 
good  oxidizing  agent.  When  the  concentrated  acid  is  heated 
with  sulfur  or  carbon  or  various  other  substances,  oxidation 
takes  place,  the  sulfuric  acid  decomposing  according  to  the 
equation 

H2S04— *H2S03 


SULFUR  AND  ITS  COMPOUNDS  155 

3.  Action  on  metals.  In  dilute  solution  sulfuric  acid  acts 
upon  many  metals,  such  as  zinc,  forming  a  sulfate  and  lib- 
erating hydrogen.  When  the  concentrated  acid  is  employed, 
the  first  action  is  one  of  oxidation.  With  copper  the  re- 
action is  represented  by  the  equation 

Cu  +  H2S04  —  >-  CuO  +  H20  +  S02 

The  copper  oxide  then  dissolves  in  an  additional  quantity 
of  sulfuric  acid  to  form  copper  sulfate  : 

CuO  +  H2SO4  —  >-  CuSO4  +  H2O 
These  two  equations  can  be  combined  into  the  form 
Cu  +  2  HS0—  ^CuS0  +  2  HO  +  SO 


4.  Action  on  salts.   We  have  repeatedly  seen  that  an  acid 
of  high  boiling  point,  heated  with  the  salt  of  some  acid  of 
lower   boiling   point,  will  drive   out  the  low-boiling   acid 
(p.  138).     The   boiling   point  of  sulfuric  acid   (338°)   is 
higher  than  that  of  almost  any  common  acid  ;   hence  it  is 
largely  used  in  the  preparation  of  other  acids. 

5.  Action  on   water.     Concentrated  sulfuric  acid  has  a 
very  great  affinity  for  water,  and  is  therefore  an  effective 
drying,  or  dehydrating,  agent.    Gases  which  have  no  chemi- 
cal action  upon  sulfuric  acid  can  be  freed  from  water  vapor 
by  bubbling  them  through  the  concentrated  acid. 

Not  only  can  sulfuric  acid  absorb  water,  but  it  will  often 
withdraw  the  elements  hydrogen  and  oxygen  from  a  compound 
containing  them,  decomposing  the  compound  and  combining 
with  the  water  so  formed.  For  this  reason  most  organic  sub- 
stances, such  as  sugar,  wood,  cotton  and  woolen  fiber,  and  even 
flesh,  all  of  which  contain  much  oxygen  and  hydrogen  in  addi- 
tion to  carbon,  are  charred,  or  burned,  by  the  action  of  the 
concentrated  acid. 


156 


FIEST  COURSE  IN  CHEMISTRY 


Salts  of  sulfuric  acid  ;  sulfates.  The  sulf ates  form  a  very 
important  class  of  salts,  and  many  of  them  have  commercial 
uses.  Copperas  (iron  sulfate),  blue  vitriol  (copper  sulfate), 
and  Epsom  salt  (magnesium  sulfate)  serve  as  examples. 
Many  sulfates  are  important  minerals,  prominent  among 
these  being  gypsum  (calcium  sul- 
fate) and  barite  (barium  sulfate). 

Monobasic  and  dibasic  acids.  Acids 
like  hydrochloric  and  nitric  acids, 
which  have  only  one  replaceable 
hydrogen  atom  in  the  molecule,  or, 
in  other  words,  which  yield  one  hy- 
drogen ion  in  solution,  are  called 
monobasic  acids.  Acids  yielding  two 
hydrogen  ions  in  solution  are  called 
dibasic  acids.  Similarly,  we  may  have 
tribasic  and  tetrabasic  acids.  The 
three  acids  of  sulfur  are  dibasic  acids. 
Normal  and  acid  salts.  It  is  pos- 
sible for  such  acids  as  H2S,  H2SO3, 
H2SO4,  to  form  two  kinds  of  salts. 
In  the  one  all  of  the  hydrogen  of  the 
acid  has  been  replaced  by  a  metal, 


FIG.  75.    A  furnace  for 
the  manufacture  of  car- 
bon disulfide 


as   in   the  salts  Na2S  and  NaaSO4. 


These  are  called  normal  salts.  In  the  other  only  one  half  of 
the  hydrogen  has  been  replaced,  as  in  the  salts  NaHS  and 
NaHSO4.  These  are  called  acid  salts,  since  they  are  at  once 
both  salts  and  acids.  Acid  salts  are  often  designated  by  the 
prefix  bi-;  thus,  NaHSO4  is  called  sodium  acid  sulfate, 
sodium  hydrogen  sulfate,  or  sodium  bisulfate. 

Carbon  disulfide  (CS2).  When  sulfur  vapor  is  passed  over 
highly  heated  carbon  the  two  elements  combine,  forming 
carbon  disulfide  (CS  ),  just  as  oxygen  and  carbon  unite  to 


SULFUR  AND  ITS  COMPOUNDS  157 

form  carbon  dioxide  (CO2).  The  substance  is  a  heavy, 
colorless  liquid,  possessing,  when  pure,  a  pleasant,  ethereal 
odor.  On  standing  for  some  time,  especially  when  exposed 
to  sunlight,  it  undergoes  a  slight  decomposition  and  ac- 
quires a  most  disagreeable  odor.  It  is  a  very  good  solvent 
for  many  substances,  such  as  gums,  resins,  waxes,  and  fats, 
which  are  insoluble  in  most  liquids.  It  boils  at  a  low  tem- 
perature (46°),  and  its  vapor  is  not  only  very  poisonous 
but  inflammable  as  well.  It  burns  in  the  air  to  form  carbon 
dioxide  and  sulfur  dioxide.  It  is  prepared  in  considerable 
quantities  for  use  as  a  solvent  and  as  an  insecticide. 

Commercially,  carbon  disulfide  is  made  by  the  direct  com- 
bination of  carbon  and  sulfur,  the  heat  necessary  for  this  union 
being  derived  from  an  electric  current.  The  main  part  of  a 
large  furnace  (Fig.  75,  A)  is  filled  with  charcoal  introduced 
through  the  trap  C.  Sulfur  is  added  through  the  hoppers  D,  D. 
An  electric  current  is  passed  in  at  E,  E.  The  heat  generated 
is  sufficient  to  vaporize  the  sulfur,  which  then  unites  with  the 
hot  carbon  to  form  carbon  disulfide.  The  vapors  escape  at  H 
and  are  condensed.  Some  of  the  furnaces  are  40  ft.  in  height 
and  yield  as  much  as  25,000  Ib.  of  the  disulfide  in  twenty-four 
hours. 

EXERCISES 

1.  Would  the  same  amount  of  sulfur  dioxide  be  formed  by  the 
combustion  of  1  g.  of  each  of  the  modifications  of  sulfur  ? 

2.  Is  the   equation  for  the   preparation   of   hydrogen   sulfide   a 
reversible  one  ?  As  ordinarily  carried  out,  does  the  reaction  complete 
itself  ? 

3.  Suppose  that    hydrogen    sulfide   were    a  liquid,  would    it  be 
necessary  to  modify  the  method  of  preparation? 

4.  Does  perfectly  dry  hydrogen  sulfide  change  the  color  of  litmus 
paper?    State  reason  for  your  answer. 

5.  What  is   an  acid    anhydride?     Aside    from  those   of  sulfur, 
what  others  have  been  mentioned  ? 


158  FIRST  COURSE  IN  CHEMISTRY 

6.  How  would  you  expect  dilute  sulfuric  acid  to  act  upon  iron? 
upon  silver?    (Refer  to  displacement  series.) 

7.  Can  you  suggest  a  reason  for  silver  spoons  becoming  tarnished 
when  in  contact  with  certain  kinds  of  food  ? 

8.  Mention  other  instances  of  catalysis  aside  from  those  given 
in  this  chapter. 

9.  In  the  commercial  preparation  of  carbon  disulfide,  what  is 
the  function  of  the  electric  current? 

10.  How  many  pounds  of  sulfur  would  be  necessary  in  the  prep- 
aration of  100  Ib.  of  98  per  cent  sulfuric  acid?  A  ns.  32.07  g. 

TOPICS  FOR  THEMES 

The  mining  of  sulfur  (see  encyclopedia). 

Goodyear  and  his  discovery  of  the  method  for  vulcanizing  rubber 
(see  encyclopedia). 

The  value  of  spraying  fruit  trees.  (Write  to  the  director  of  your 
experiment  station  for  information.) 


CHAPTER  XVIII 
THE  PERIODIC  LAW 

A  number  of  the  elements  have  now  been  studied  some- 
what closely.  The  first  three  of  these,  oxygen,  hydrogen, 
and  nitrogen,  while  having  some  properties  in  common  with 
each  other,  have  almost  no  point  of  similarity  as  regards 
their  chemical  conduct.  On  the  other  hand,  oxygen  and 
sulfur,  while  quite  different  physically,  have  much  in  com- 
mon in  the  way  they  act  toward  other  chemicals. 

About  eighty  elements  are  now  known.  If  all  of  these 
should  have  properties  as  diverse  as  do  oxygen,  hydrogen, 
and  nitrogen,  the  study  of  chemistry  would  plainly  be  very 
difficult  and  complicated.  Fortunately  a  study  of  the  ele- 
ments shows  that  certain  ones  resemble  each  other  more 
or  less  closely,  so  that  it  is  possible  to  divide  them  into 
groups  and  then  study  the  group  as  a  whole.  A  number  of 
different  methods  of  classifying  the  elements  have  been 
suggested,  but  that  advanced  in  1869  by  the  Russian  chem- 
ist Mendeleeff  (Fig.  76)  has  proved  the  most  fruitful.  In 
accordance  with  this  method,  the  elements  are  arranged  in 
groups  or  periods,  according  to  their  atomic  weights. 

The  periodic  grouping.  The  general  arrangement  sug- 
gested by  Mendeleeff  and  extended  so  as  to  include 
elements  more  recently  discovered  is  as  follows:  Omit- 
ting hydrogen,  which  has  the  smallest  atomic  weight,  and 
beginning  with  helium,  which  has  an  atomic  weight  of  3.99, 
the  succeeding  seven  elements  are  arranged  in  a  horizontal 

159 


160 


FIKST  COUKSE  IX  CHEMISTRY 


row  in  the  order  of  their  atomic  weights,  as  given  below. 
These  eight  elements  all  differ  markedly  from  each  other, 
but  the  one  having  the  next  highest  atomic  weight,  neon,  is 
very  similar  to  helium.  It  is  placed  just  under  helium, 
and  a  new  horizontal  row  follows  as  shown  below.  The 
next  element,  argon,  resembles  helium  and  neon  and  begins 
a  third  row. 

He  (3.99)  Li  (6.94)  Gl  (9.1)       B  (11)      C  (12)      N  (14.01)  O  (16)        F  (19) 
Ne  (20.2)  Na  (23)     Mg  (24.32)  Al  (27.1)  Si  (28.3)  P  (04)       S  (32.07)   Cl  (35.46) 
A  (39.88)   K  (39.1)  Ca  (40.07)  Sc  (44.1)  Ti  (48.1)  V  (51)      Cr  (52)      Mn  (54.93) 

If  now  we  consider  the  elements  that  fall  under  each 
other  in  these  three  rows,  a  remarkable  fact  is  brought 

to  light.  Not  only  is  there 
a  strong  similarity  between 
helium,  neon,  and  argon, 
which  form  the  first  vertical 
column,  but  the  elements  in 
the  other  columns  exhibit 
much  of  the  same  kind  of 
similarity  among  themselves, 
and  evidently  form  natural 
groups. 

Iron,  nickel,  and  cobalt, 
following  manganese,  have 
atomic  weights  near  together, 
and  are  very  similar  chemi- 
cally. They  do  not  strongly 


FIG.  76.   Mendel^eff  (1834-1907) 

A  famous  Russian  chemist,  who  first 
proposed  the  periodic  classification  of 


the  elements 


resemble  any  of  the  elements 
so  far  considered,  and  so  are 
placed  in  a  group  by  them- 
selves. The  first  three  horizontal  rows  of  the  table  (p.  161) 
shows  the  arrangement  of  these  twenty-seven  elements. 
A  new  horizontal  row  is  begun  with  copper. 


THE  PERIODIC  LAW 


161 


SS2 


*s 


Is 


!! 

"o  o 

e3    c8 


162  FIKST  COURSE  IN  CHEMISTRY 

Following  the  fifth  and  seventh  rows  are  groups  of  three 
closely  related  elements,  so  the  completed  arrangement  has 
the  appearance  represented  in  the  table. 

The  relation  of  properties  of  elements  to  atomic  weights. 
There  is  evidently  a  close  relation  between  the  properties  of 
an  element  and  its  atomic  weight.  For  example,  consider 
the  elements  in  the  first  horizontal  row.  Helium  is  an  inert 
element.  Following  it,  lithium  is  a  metallic  element,  has  a 
valence  of  1,  and  possesses  a  strong  base-forming  character. 
The  next  element,  glucinum,  has  a  valence  of  2,  and  is  less 
strongly  base-forming,  while  boron  has  some  base-forming 
and  some  acid-forming  properties.  In  carbon,  all  base- 
forming  properties  have  disappeared,  and  the  acid-forming 
properties  are  more  marked  than  in  boron.  These  become 
still  more  emphasized  as  we  pass  through  nitrogen  and 
oxygen,  until  on  reaching  fluorine  we  have  one  of  the 
strongest  acid-forming  elements.  The  properties  of  these 
eight  elements  vary  regularly  with  their  atomic  weights, 
or,  in  mathematical  language,  are  regular  functions  of  them. 

The  periodic  law.  If  it  were  true  that  helium  had  the 
smallest  atomic  weight  of  any  of  the  elements  and  fluorine 
the  greatest,  so  that  in  passing  from  one  to  the  other  we 
included  all  the  elements,  we  could  say  that  the  properties  of 
elements  were  continuous  functions  of  their  atomic  weights. 
But  fluorine  is  an  element  of  relatively  small  atomic  weight, 
and  the  one  following  it,  neon,  breaks  the  regular  order, 
for  in  it  reappear  all  the  characteristic  properties  of  helium. 
Sodium,  following  neon,  bears  much  the  same  relation  to 
lithium  that  neon  does  to  helium,  and  the  properties  of  the 
elements  in  the  second  row  vary  much  as  the  properties  of 
the  elements  in  the  first  row  did,  until  argon  is  reached, 
when  another  repetition  begins.  The  properties  of  the  ele- 
ments do  not  vary  continuously,  therefore,  with  atomic 


THE  PEBIODIC  LAW  163 

weights,  but  at  regular  intervals  there  is  a  repetition,  or 
period.  This  generalization  is  known  as  the  periodic  law, 
and  may  be  stated  thus :  The  properties  of  elements  are 
periodic  functions  of  their  atomic  weights. 

Two  families  in  a  group.  The  elements  of  each  group 
(excepting  Group  0)  falls  naturally  into  two  families.  The 
elements  in  the  odd-numbered  horizontal  rows,  or  periods, 
form  one  family,  those  in  the  even-numbered  periods,  the 
other.  In  the  table  these  are  arranged  under  the  headings  A 
and  B.  The  elements  in  one  family  are  much  more  similar  to 
each  other  than  they  are  to  those  in  the  other  family  in  the 
same  group.  Thus,  magnesium,  zinc,  cadmium,  and  mercury 
form  one  family  of  very  similar  elements  in  Group  II,  while 
calcium,  strontium,  barium,  and  radium  form  the  other. 

Family  resemblances.  Let  us  inquire  more  closely  in 
what  respects  the  elements  of  a  family  resemble  each  other. 

1.  Valence.    In  general  the  valence  of  the  elements  in  a 
family  is  the  same,  and  the  formulas  of  their  compounds 
are  therefore  similar.    The  formulas  R2O,  RO,  etc.,  placed 
below  the  columns,  represent  the  oxides  of  the  elements  in 
the  column,  while  the  formulas  RH,  RH2,  etc.  represent 
the  hydrides  or  chlorides. 

2.  Chemical  conduct.    The  chemical  characteristics  of  the 
members  of  a  family  are  somewhat  similar.    If  one  member 
is  a  metal,  the  others  usually  are ;  if  one  is  a  nonmetal,  so, 
too,   are  the   others.     There   is    also   a  certain  regularity 
in  the  properties  of  the  elements  in  each  family.    If  the  one 
element  at  the  head  of  the  family  is  a  strong  acid-forming 
element,  this  characteristic  is  likely  to  diminish  gradually 
as  we  pass  to  the  members  of  the  family  having  higher 
atomic  weights.   Thus,  phosphorus  is  strongly  acid-forming, 
arsenic  less  so,  and  antimony  still  less  so,  while  bismuth 
has  almost  no.  acid-forming  properties. 


164 


FIRST  COURSE 


CHEMISTRY 


3.  Physical  properties.  In  the  same  way,  the  physical 
properties  of  the  members  of  a  family  are  in  general  some- 
what similar,  and  show  a  regular  gradation  as  we  pass  from 
element  to  element  in  the  family.  Thus,  the  densities  of  the 
members  of  the  magnesium  family  increase  with  their  atomic 
weights,  while  their  melting  points  decrease  (p.  313). 

Value  of  the  periodic  law.  The  periodic  law  has  proved  of 
much  value  in  the  development  of  the  science  of  chemistry. 

1.  It  simplifies  study.    It  is  at  once  evident  that  such 
regularities   very  much  simplify  the  study  of   chemistry. 
A  thorough  study  of  one  element  of  a  family  makes  the 
study  of  the  other  members  a  much  easier  task. 

2.  It  suggests  the  probable  existence  of  new  elements.  When 
the  periodic  law  was  first  formulated  there  were  a  number 
of  vacant  places  in  the  table  which  evidently  belonged  to 
elements  at  that  time  unknown.    From  their  position  in 
the  table,  Mendeleeff  predicted  with  great  precision  the 
properties  of  the  elements  which  he  felt  sure  would  one  day 
be  discovered  to  fill  these  places.    Three  of  them,  scandium, 
germanium,  and  gallium,  were  found  within  fifteen  years, 
and  their  properties  agreed  in  a  remarkable  way  with  the 
predictions  of  Mendeleeff. 

This  is  shown  in  the  following  table,  in  which  the  proper- 
ties of  gallium  are  compared  with  those  which  Mendeleeff 
predicted : 


PROPERTIES  OF  GALLIUM 

PREDICTED 

FOUND 

Atomic  weight    
Melting  point     

about  69 
low 

69.9 
30.2° 

Specific  gravity  

5.9 

5.95 

Formula  of  oxide    

X  O 

Ga  O 

Action  of  air 

^2^3 

no  action 

J.  only  slight,  even 

\      at  red  heat 

THE  PERIODIC  LAW  165 

3.  It  indicates  probable  errors.  The  physical  constants  of 
many  of  the  elements  did  not  at  first  agree  with  those  sug- 
gested by  the  periodic  law,  and  a  further  study  of  many 
such  cases  showed  that  errors  had  been  made. 

Imperfections  of  the  law.  There  still  remain  a  good  many 
features  which  must  be  regarded  as  imperfections  in  the 
law.  Most  conspicuous  is  the  fact  that  the  element  hydro- 
gen has  no  place  in  the  table.  Moreover,  according  to  their 
atomic  weights,  tellurium  should  follow  iodine,  and  argon 
should  follow  potassium,  but  their  properties  show  in  each 
case  that  this  order  must  be  reversed.  The  table  separates 
some  elements  altogether  which  in  many  respects  have 
closely  agreeing  properties.  Iron,  chromium,  and  manga- 
nese, although  they  are  similar  in  many  respects,  are  all  in 
different  groups. 

The  periodic  law  is  therefore  to  be  regarded  as  but  a 
partial  and  imperfect  expression  of  some  very  important 
and  fundamental  relation  between  the  substances  which  we 
know  as  elements,  the  exact  nature  of  this  relation  being 
as  yet  not  completely  clear  to  us. 

EXERCISES 

1.  Suppose  that  an  element  were  discovered  that  filled  the  blank 
in  Group  0,  Period  4  ;  what  properties  would  it  probably  have  ? 

2.  Sulfur  and  oxygen  both  belong  to  the  same  group;   in  what 
respects  are  they  similar? 

TOPIC  FOR  THEMES 

Mendeleeft'  (Thorpe,  Essays  in  Historical  Chemistry). 


CHAPTER  XIX 
THE  CHLORINE  FAMILY 


NAME 

ATOMIC 
WEIGHT 

MELTING 
POINT 

BOILING 
POINT 

COLOH  AND  STATE 

Fluorine  (F)     .     . 

19.00 

-223° 

-187° 

Pale-yellowish  gas 

Chlorine  (Cl)    .     . 

35.46 

-101.5° 

-  33.6° 

Greenish-yellow  gas 

Bromine  (Br)    . 

79.92 

-  7.3° 

63° 

Red  liquid 

Iodine  (I) 

126.92 

113.5° 

184.4° 

Purplish-black  solid 

The  family.  The  four  elements  named  in  the  above  table 
form  a  strongly  marked  family,  and  illustrate  very  clearly 
the  way  in  which  the  members  of  a  family  in  a  periodic 
group  resemble  each  other,  as  well  as  the  character  of  the 
differences  which  we  may  expect  to  find  between  the  in- 
dividual members.  The  elements  constituting  the  family 
are  often  termed  the  halogens.  They  will  be  described  in 
the  order  of  their  atomic  weights. 


FLUORINE 

Occurrence.  Fluorine  occurs  in  nature  most  abundantly 
in  the  mineral  fluorite  (CaF2),  in  cryolite  (Na3AlF6),  and 
in  the  complex  mineral  fluorapatite  (3  Ca3(PO4)2  •  CaF2). 

Preparation.  Because  of  its  great  activity,  the  element 
is  difficult  to  liberate  from  its  compounds.  Its  prepara- 
tion was  finally  accomplished  in  1886  by  the  French 
chemist  Moissan  (Fig.  77)  by  the  electrolysis  of  hydrogen 
fluoride.  His  success  was  clue  to  his  observation  that 

166 


THE  CHLORINE  FAMILY 


167 


liquid  hydrogen  fluoride  (H2F2)  in  which  is  dissolved  a 
little  potassium  acid  fluoride  (KHF2)  is  a  good  electrolyte. 
The  electrolyte  was  placed  in  a  U-shaped  tube  made  of  plati- 
num (or  copper)  upon  which  fluorine  has  little  action.  This 
tube  was  furnished  with  electrodes  and  delivery  tubes  as 
shown  in  Fig.  78.  When  the  solution  was  electrolyzed,  hy- 
drogen was  set  free  at  the  cathode,  and  fluorine  at  the  anode. 


rS!J&Sfr;JT 


r .  r"v  A-  £>£••-  -^  -^ ' 


FIG.  77.    Tablet  erected  by  the  associates  and  friends  of  Moissan,  in  his 
laboratory  in  Paris,  in  1906,  on  the  twentieth  anniversary  of  the  isolation 

of  fluorine 


Properties  and  conduct.  Fluorine  is  a  gas  of  slightly 
yellowish  color  and  is  very  difficult  to  liquefy.  It  is  the 
most  active  of  all  the  elements  at  ordinary  temperatures, 
and  combines  with  all  the  common  elements  save  oxygen, 
very  often  with  incandescence.  It  has  a  particularly  strong 
affinity  for  hydrogen,  and  is  able  to  withdraw  it  from  any 
of  its  compounds  with  other  elements,  forming  hydrogen 
fluoride.  It  is  extremely  poisonous. 


168 


FIRST  COURSE 


CHEMISTRY 


Hydrogen  fluoride  (H2F2).  Hydrogen  fluoride  is  readily 
obtained  from  fluorite  by  the  action  of  concentrated  sul- 
furic  acid.  The  equation  is  as  follows  : 


CaF2  +  H2SO4 


CaSO 


HaFa 


In  its  properties  hydrogen  fluoride  resembles  the  hydrides 
of  the  other  elements  of  this  family,  being,  however,  more 
easily  condensed  to  a  liquid.  It 
boils  at  19.4°,  and  can  therefore 
be  liquefied  at  ordinary  pressures. 
It  is  soluble  in  all  proportions  in 
water,  forming  hydrofluoric  acid.  Its 
fumes  are  exceedingly  irritating  to 
the  respiratory  organs,  and  several 
chemists  have  lost  their  lives  by 
accidentally  breathing  them. 

Hydrofluoric  acid.  Hydrofluoric 
acid,  like  other  strong  acids,  readily 
acts  on  bases  and  metallic  oxides, 
and  forms  the  corresponding  fluor- 
ides. It  acts  very  vigorously  upon 
organic  matter,  a  single  drop  of  the  concentrated  acid 
making  a  sore  on  the  skin  which  is  slow  in  healing  and  very 
painful.  Its  most  characteristic  property  is  its  action  upon 
silicon  dioxide  (SiO2),  with  which  it  forms  w^ater  and  the 
gas  silicon  tetrafluoride  (SiF4),  as  shown  in  the  equation 


FIG.  78.   Apparatus  for  the 
preparation  of  fluorine 


SiOa+2HaFs 


SiF4+2H20 


Glass  consists  of  certain  compounds  of  silicon  which  are 
likewise  acted  on  by  the  acid,  so  that  it  cannot  be  kept  in 
glass  bottles.  It  is  preserved  in  flasks  made  of  a  wax  derived 
from  petroleum  and  known  as  ceresin. 


THE  CHLOEINE  FAMILY  169 

Etching.  Advantage  is  taken  of  this  reaction  in  etching 
designs  upon  glass.  The  glass  vessel  is  painted  over  with  a 
protective  paint  upon  which  the  acid  will  not  act,  the  parts 
which  it  is  desired  to  make  opaque  being  left  unprotected.  A 
mixture  of  fluorite  and  sulfuric  acid  is  then  painted  over  the 
vessel,  and  after  a  few  minutes  the  vessel  is  washed  clean. 
AYherever  the  hydrofluoric  acid  comes  in  contact  with  the  glass 
it  acts  upon  it,  destroying  its  luster  and  making  it  opaque,  so 
that  the  exposed  design  will  be  etched  upon  the  clear  glass. 
Frosted  glass  globes  are  often  made  in  this  way,  but  more 
frequently  by  a  sand  blast. 

The  etching  may  also  be  effected  by  covering  the  glass  with 
a  thin  layer  of  paraffin,  cutting  the  design  through  the  wax 
and  then  exposing  the  glass  to  the  fumes  of  the  gas. 

CHLORINE 

Historical.  While  studying  the  action  of  hydrochloric 
acid  upon  the  mineral  pyrolusite,  the  Swedish  chemist, 
Scheele,  in  1774,  obtained  a  yellowish  gaseous  substance 
to  which  he  gave  a  name  in  keeping  with  the  phlogiston 
theory  then  current  (p.  6).  Later  it  was  supposed  to  be 
a  compound  containing  oxygen.  In  1810,  however,  the 
English  chemist,  Sir  Humphry  Davy,  proved  it  to  be  an 
element  and  named  it  chlorine. 

Occurrence.  Chlorine  does  not  occur  free  in  nature,  but 
its  compounds  are  widely  distributed.  It  is  found  in  com- 
bination with  the  metals  in  the  form  of  chlorides,  those  of 
sodium,  magnesium,  and  potassium  being  the  most  abun- 
dant. All  salt  water  contains  these  substances,  particularly 
sodium  chloride  (common  salt),  and  very  large  salt  beds  of 
chlorides  are  found  in  many  parts  of  the  world. 

Preparation.  Two  general  methods  of  preparing  chlorine 
may  be  mentioned,  one  of  which  is  the  usual  laboratory 
method,  while  the  other  is  the  commercial  method. 


170 


FIRST  COUESE  IN  CHEMISTRY 


1.  Laboratory  method.  In  the  laboratory,  chlorine  is  easily 
made  by  warming  the  mineral  pyrolusite  (manganese 
dioxide,  MnO2)  with  concentrated  hydrochloric  acid.  The 
first  reaction,  which  is  similar  to  the  action  of  acids  upon 
oxides  in  general,  is  expressed  in  the  equation 

MnO2  +  4  HC1  --  *  MnCl4  -f  2  H2O 

The  manganese  compound  so  formed  is  very  unstable,  how- 
ever, and  breaks  down  according  to  the  equation 


MnCl 


MnCl  +  C1 


Instead  of  using  hydrochloric  acid  in  the  preparation  of 
chlorine,  it  serves  just  as  well  to  use  a  mixture  of  sodium 
chloride  and  sulfuric  acid,  since  these  two  react  to  form 
hydrochloric  acid.  In  this  case  the  complete  reaction  is 
expressed  by  the  equation 

2  NaCl+MnOa+2  H2SCT- 


O 


•Na2S04+MnS04+2  H2O+C12 

The  manganese 
dioxide  and  the  hy- 
drochloric acid  are 
brought  together  in 
a  flask  A  (Fig.  79), 
and  a  gentle  heat  is 
applied.  The  chlo- 
rine set  free  passes 
out  through  B,  bub- 
bles through  a  little 
water  in  C  (which 
removes  any  hydro- 
gen chloride  carried 
over  with  it),  and 

finally  through  some  sulfuric  acid  in  D,  which  dries  the  gas. 
Being  somewhat  soluble  in  water,  it  is  collected  in  cylinder  E  by 
displacement  of  air,  the  color  showing  when  the  cylinder  is  full. 


FIG.  79.   The  preparation  of  pure  chlorine 


THE  CHLOEINE  FAMILY  171 

2.  Commercial  method.  It  will  be  recalled  that  when  a 
solution  of  sodium  chloride  is  electrolyzed,  chlorine  is 
evolved  at  the  anode,  while  sodium  hydroxide  is  formed  at 
the  cathode  and  remains  in  solution  (p.  104).  All  of  the 
chlorine  prepared  for  commercial  purposes  in  the  United 
States  is  obtained  in  this  way.  The  method  has  advan- 
tages in  that  sodium  chloride  is  cheap  and  that  the  sodium 
hydroxide  formed  in  the  process  has  many  commercial  uses. 
The  chlorine  so  obtained  is  either  compressed  in  strong  iron 
cylinders  and  shipped  in  this  form  or  is  passed  into  slaked 
lime  forming  the  solid  known  as  chloride  of  lime  or  bleach- 
ing powder  (p.  303),  which  can  be  easily  shipped  and  from 
which  the  chlorine  can  be  recovered  as  needed. 

Properties.  Chlorine  is  a  greenish-yellow  gas  which  has  a 
peculiar,  suffocating  odor  and  produces  a  very  irritating  effect 
upon  the  throat  and  lungs.  Even  when  inhaled  in  small  quan- 
tities, it  often  produces  all  the  symptoms  of  a  very  hard  cold, 
and  in  larger  quantities  may  have  serious  and  even  fatal 
effects.  Inhaling  ether  or  ammonia  gives  some  relief.  Chlo- 
rine is  nearly  2.5  times  as  heavy  as  air,  and  can  therefore  be 
collected  by  displacement  of  air.  One  volume  of  water  under 
ordinary  conditions  dissolves  about  3  volumes  of  chlorine. 

Chemical  conduct.  At  ordinary  temperatures  chlorine  is 
far  more  active  chemically  than  any  of  the  elements  we 
have  so  far  considered,  with  the  exception  of  fluorine ; 
indeed,  it  is  one  of  the  most  active  of  all  elements. 

1.  Action  on  elements.  A  great  many  elements  combine 
directly  with  chlorine,  especially  when  hot.  A  strip  of  cop- 
per foil  heated  in  a  burner  flame  and  then  dropped  into  chlo- 
rine burns  with  incandescence.  Antimony  and  arsenic  in  the 
form  of  a  fine  powder  at  once  burst  into  flame  when  dropped 
into  jars  of  chlorine.  The  products  formed  in  all  cases  where 
chlorine  combines  with  another  element  are  called  chlorides. 


172 


FIRST  COURSE  IN  CHEMISTEY 


2.  Action  on  hydrogen.    Chlorine  has  a  particularly  strong 

affinity  for   hydrogen,  uniting  with  it  to  form  hydrogen 

chloride.  A  jet  of  hydrogen 
burning  in  the  air  continues 
to  burn  when  introduced 
into  a  jar  of  chlorine  A 
(Fig.  80),  giving  a  some- 
what luminous  flame.  A 
mixture  of  the  two  gases 
explodes  violently  when  a 
spark  is  passed  through  it 
or  when  it  is  exposed  to 
bright  sunlight. 

3.  Action  on  substances 
containing  hydrogen.  Not 
only  will  chlorine  combine 
directly  with  free  hydrogen 

but  it  will  often  abstract  the  element  from  its  compounds. 

Thus,  when  chlorine  is  passed  into  a 

solution    containing    hydrogen    sulfide, 

sulfur  is  precipitated  and  hydrochloric 

acid  formed.    The  reaction  is  shown  by 

the  following  equation : 


FIG.  80.    Burning  hydrogen  in 
chlorine 


Even  water,  which  is  a  very  stable 
compound,  can  be  decomposed  by  chlo- 
rine, the  oxygen  being  liberated.  This 
may  be  shown  in  the  following  way: 


FIG.  81.  The  action  of 
chlorine  on  water  in 
the  sunlight 


Action  of  chlorine  on  water.   A  long  tube 
of  rather  large  diameter  is  filled  with  a 

concentrated  solution  of  chlorine  in  water  and  inverted  in  a 
vessel   of  the   same   solution,   as   shown   in   Fig.  81,  and  the 


THE  CHLORINE  FAMILY 


173 


apparatus  is  placed  in  bright  sunlight.  Very  soon  bubbles  of  a 
gas  will  be  observed  to  rise  through  the  solution  and  collect  in 
the  tube,  and  an  examination  of  this  gas  will  show  that  it  is  oxy- 
gen. It  is  liberated  from  water  in  accordance  with  the  following 
equation  :  2  ^  +  2  ^  -  ^  4  HC1  +  ^ 

4.  Action  on  color  substances  ;  bleaching  action.  If  strips  of 
brightly  colored  cloth  or  some  highly  colored  flowers  are 
placed  in  dry  chlorine,  no 
marked  change  in  color  is 
noticed,  as  a  rule.  If,  how- 
ever, the  cloth  and  flowers 
are  first  moistened,  the  color 
rapidly  disappears,  or,  in 
other  words,  the  objects 
are  bleached.  Evidently  the 
moisture  as  well  as  the 
chlorine  is  concerned  in 
the  action.  A  study  of  the 
case  shows  that  the  chlo- 
rine combines  with  the  hy- 
drogen of  the  water,  and 
the  oxygen  set  free  oxi- 


EIG< 


The  bleaching  action  of 
moist  chlorine 


dizes  the   color  substance, 

converting  it  into  a  color- 

less compound.    It  is  evident  from  this  explanation  that 

chlorine  will  bleach  only  those  substances  which  are  changed 

into  colorless  compounds  by  oxidation. 

Fig.  82  illustrates  the  bleaching  action  of  chlorine.  Strips 
from  the  same  piece  of  cloth  are  suspended  in  three  jars, 
of  which  the  first  contains  air,  the  second  dry  chlorine,  and 
the  third  moist  chlorine.  It  will  be  noted  that  dry  chlorine 
has  almost  no  bleaching  action,  while  the  moist  chlorine  has 
partially  removed  the  color. 


1T4 


FIRST  COURSE  IN  CHEMISTRY 


5.  Action  as  a  disinfectant.  Chlorine  has  also  marked 
germicidal  properties,  and  the  free  element,  as  well  as  com- 
pounds from  which  it  is  easily  liberated,  are  used  as  disin- 
fectants. It  is  also  used  to  destroy  the  microorganisms  in 
city  water  supplies. 

Bleaching.  The  process  known  as  bleaching  is  an  im- 
portant one  in  connection  with  many  industries.  Thus, 

the  various  kinds  of  fab- 
rics woven  from  vegetable 
fibers,  such  as  flax  and 
cotton,  are  always  more  or 
less  colored  ;  hence  bleach- 
ing is  necessary  if  a  white 
fabric  is  desired.  This  was 
formerly  accomplished  by 
spreading  the  cloth  on  plots 
of  grass  and  exposing  it  to 
the  action  of  air  and  sun- 
light, but  the  process  was 


very  slow.   The  same  results 

are  now  obtained  in  a  short 

time  by  the  use  of  chlorine. 

Certain    foods,   such   as 


FIG.  83.    Laboratory  preparation  of 
hydrogen  chloride 


dried  fruits  (p.  150)  and  the  lower  grades  of  flour,  are  also 
bleached.  The  bleaching  of  flour  was  formerly  prohibited, 
but  is  now  largely  practiced,  since  it  makes  the  product  white. 

Other  bleaching  agents.  While  chlorine  is  the  bleaching  agent 
most  generally  used,  a  number  of  other  substances,  such  as 
sulfurous  acid,  hydrogen  dioxide,  nitrogen  dioxide,  and  ozone, 
are  sometimes  employed.  Chlorine  injures  some  fabrics,  such 
as  silk  and  straw  goods,  and  in  such  cases  sulfurous  acid  is 
used.  Flour  is  bleached  either  by  chlorine  or  by  nitrogen 
dioxide,  a  very  small  amount  sufficing  to  remove  the  color. 


THE  CHLORINE  FAMILY  175 

Nascent  state.  It  will  be  noticed  that  when  oxygen  is 
set  free  from  water  by  chlorine,  it  is  at  that  instant  able 
to  do  what  ordinary  oxygen  gas  cannot  do,  for  it  bleaches 
substances  which  would  remain  unchanged  in  dry  air  or 
pure  oxygen.  It  is  generally  true  that  the  activity  of  an 
element  is  greatest  at  the  instant  of  its  liberation  from  its 
compounds.  To  express  this  fact,  elements  at  the  instant  of 
liberation  are  said  to  be  in  the  nascent  state.  When  moist 
chlorine  acts  as  a  bleaching  agent,  it  is  nascent  oxygen 
which  does  the  bleaching. 

Hydrogen  chloride  (HC1).  The  preparation  of  this  gas 
may  be  discussed  under  two  general  heads: 

1.  Laboratory  preparation.   While  hydrogen  chloride  (in 
solution  known  as  hydrochloric  acid)  can  be  prepared  by 
burning  hydrogen  in  chlorine,  it  is  much  more  conveniently 
obtained  by  treating  common  salt  (sodium  chloride)  with 
sulfuric  acid.    The  following  equation  shows  the  reaction : 

2  NaCl  -f  H2SO4 >•  Na2SO4  +  2  HC1 

The  dry  salt  is  placed  in  the  flask  A  (Fig.  83),  sulfuric  acid 
is  added,  and  the  flask  gently  warmed.  The  hydrogen  chloride 
is  rapidly  given  off,  and  can  be  collected  by  displacement  of 
air.  To  prepare  a  solution  of  the  gas,  the  end  of  the  delivery 
tube  is  fixed  just  above  the  level  of  some  water  in  the  cylin- 
der B.  The  gas  is  very  soluble,  and  is  absorbed  as  fast  as  it 
escapes  from  the  tube. 

2.  Commercial  preparation.    Commercially,  hydrogen  chlo- 
ride  is  prepared  in  connection  with  the  manufacture  of 
sodium  sulfate,  the  reaction  being  the  same  as  that  just 
given.   It  is  also  prepared  by  heating  sodium  hydrogen  sul- 
fate (which  is  obtained  in  the  manufacture  of  nitric  acid) 
(p.  128)  and  sodium  chloride: 

NaCl  4-  NaHSO4 >-  NaoSO4  4-  HC1 


176 


FIRST  COUESE  IN  CHEMISTEY 


In  either  case  the  hydrogen  chloride  liberated  is  passed 
into  water,  in  which  it  dissolves,  the  solution  forming  the 
hydrochloric  acid  of  commerce.  When  the  materials  are 
pure,  a  colorless  solution  is  obtained.  The  most  concen- 
trated solution  has  a  density  of  about  1.2  and  contains 
approximately  40  per  cent  of  hydrogen  chloride.  The  com- 
mercial acid  (often  called  muri- 
atic acid)  is  usually  colored  yellow 
by  impurities. 

Composition  of  hydrogen  chlo- 
ride. When  measured  volumes 
of  hydrogen  and  chlorine  are 
caused  to  unite,  it  is  found  that 
1  volume  of  hydrogen  combines  with 

1  volume  of  chlorine.   Other  experi- 
ments show  that   the  volume  of 
hydrogen  chloride  formed  is  just 
equal  to  the  sum  of  the  volumes 
of  hydrogen  and  chlorine.   There- 
fore 1  volume  of  hydrogen  combines 
with  1  volume  of  chlorine  to  form 

2  volumes    of   hydrogen    chloride. 
Since  chlorine  is  35.24  times  as 

heavy  as  hydrogen,  it  follows  that  1  part  of  hydrogen 
by  weight  combines  with  35.24  parts  of  chlorine  to  form 
36.24  parts  of  hydrogen  chloride. 

Properties.  Hydrogen  chloride  is  a  colorless  gas  which 
has  an  irritating  effect  when  inhaled,  but  no  marked  odor. 
It  is  heavier  than  air  and  is  very  soluble  in  water.  Under 
standard  conditions,  1  volume  of  water  dissolves  about  500 
volumes  of  the  gas.  On  warming  such  a  solution,  the  gas 
escapes  until,  at  the  boiling  point,  the  solution  contains 
about  20  per  cent  by  weight  of  HC1.  Further  boiling  will 


FIG.  84.    The  hydrogen  chlo- 
ride fountain  illustrating  the 
solubility  of  the  gas 


THE  CHLOEINE  FAMILY  177 

not  drive  out  any  more  acid,  but  the  solution  will  distill 
with  unchanged  concentration.  A  more  dilute  solution  than 
this  will  lose  water  on  boiling  until  it  has  reached  the  same 
concentration,  20  per  cent,  and  will  then  distill  unchanged. 

The  extreme  solubility  of  hydrogen  chloride  in  water  may 
be  shown  as  follows  :  A  perfectly  dry  flask  A- (Fig.  84)  is  filled 
with  hydrogen  chloride.  This  flask  is  connected,  by  means  of 
a  glass  tube,  with  a  similar  flask  B,  which  is  nearly  filled  with 
water,  as  shown  in  the  figure.  The  end  of  the  tube  opening 
into  flask  A  is  drawn  out  to  a  rather  fine  jet.  By  blowing  into 
the  tube  C,  a  few  drops  of  water  are  forced  into  A.  Some  of 
the  hydrogen  chloride  at  once  dissolves,  thus  diminishing  the 
pressure  inside  the  flask.  The  water  then  flows  continuously 
from  B  into  A,  until  nearly  all  the  hydrogen  chloride  is  absorbed. 
It  is  evident  that  the  connection  must  be  air-tight.  The  experi- 
ment is  more  striking  if  the  water  in  B  is  first  colored  deep 
blue  with  litmus. 

Chemical  conduct.  The  most  important  chemical  char- 
acteristics of  hydrogen  chloride  are  the  following: 

1.  Action  as  an  acid.    Its  aqueous  solution  (hydrochloric 
acid)  is  a  very  strong  acid ;  indeed,  it  is  one  of  the  strong- 
est acids  known,  and  has  many  commercial  uses.   This  acid 
acts  upon  oxides  and  hydroxides,  converting  them  into  salts: 

CuO  +  2  HC1  — *  CuCl2  +  H2O 
NaOH  +  HC1 *  NaCl  +  H2O 

It  acts  upon  many  metals,  forming  chlorides  and  liberating 

hydrogen  :          Zn  +  2  HC1 ^  ZnCl2  +  H2 

2  Al  +  6  HC1 *•  2  A1C18  +  3  H2 

2.  Action  on  oxidizing  agents.    Many  oxidizing  agents  act 
upon    hydrogen    chloride    as    expressed    in    the   following 
equation :  4  HC1  +  Q^ ^  2 II2O  +  2  C12 

The  hydrogen  combines  with  oxygen,  liberating  chlorine. 


178  FIBST  COURSE  IN  CHEMISTRY 

Aqua  regia.  Since  nitric  acid  is  a  good  oxidizing  agent, 
we  might  expect  it  to  liberate  chlorine  from  hydrogen 
chloride,  and  this  is  found  to  be  the  case.  A  mixture  of 
1  part  of  nitric  acid  and  3  parts  of  hydrochloric  acid 
is  called  aqua  regia,  and  is  one  of  the  strongest  solvents 
known.  It  owes  its  solvent  powers  not  to  its  acid  proper- 
ties but  to  the  nascent  chlorine  which  it  liberates.  Metals 
such  as  gold  and  platinum,  which  are  not  soluble  in  any 
of  the  common  acids,  readily  dissolve  in  aqua  regia,  being 
converted  into  chlorides  by  the  nascent  chlorine. 

Salts  of  hydrochloric  acid ;  chlorides.  The  chlorides  of 
all  the  metals  are  known,  and  many  of  them,  such  as  sodium 
chloride,  are  very  important  compounds.  With  very  few 
exceptions  they  are  solids,  and  all  are  soluble  in  water  ex- 
cept silver  chloride,  lead  chloride,  and  mercurous  chloride. 

Compounds  of  chlorine  with  oxygen  and  hydrogen.  Chlorine 
combines  with  oxygen  and  hydrogen  to  form  four  different 
acids.  They  are  all  quite  unstable,  and  most  of  them  cannot 
be  prepared  in  pure  form.  Their  salts  can  easily  be  made,  how- 
ever, and  some  of  them  will  be  met  with  in  the  study  of  the 
metals.  The  formulas  and  names  of  these  acids  are  as  follows : 

HC1O hypochlorous  acid 

HC1O0 chlorous  acid 

HC1O3 chloric  acid 

IIC1O4 perchloric  acid 

BKOMINE 

Historical.  Bromine  was  discovered  in  1826  by  the 
French  chemist  Ballard,  who  isolated  it  from  sea  salt.  He 
named  it  bromine  ("  stench")  because  of  its  bad  odor. 

Occurrence.  Bromine  occurs  almost  entirely  in  the  form  of 
bromides,  especially  as  sodium  bromide  and  magnesium  bro- 
mide, which  are  found  in  many  salt  springs  and  salt  deposits. 


THE  CHLOKINE  FAMILY 


179 


The  Stassfurt  deposits  in  Germany  (p.  286)  and  the  salt 
waters  of  Ohio  and  Michigan  are  especially  rich  in  bromides. 

Preparation  of  bromine.    The  laboratory  method  of  pre- 
paring bromine  is  essentially  different  from  the  commercial. 

1.  Laboratory  method.  As  in  the  case  of  chlorine,  bromine 
can  be  prepared  by  the  action  of  hydrobromic  acid  (HBr) 
on  manganese  dioxide. 
Since  hydrobromic  acid 
is  not  a  common  article 
of  commerce,  a  mixture 
of  snlfuric  acid  and  a 
bromide  is  commonly 
substituted  for  it.  The 
materials  are  placed  in 
a  retort  A,  arranged  as 
shown  in  Fig.  85.  The 
end  of  the  retort  comes 
close  to  the  surface  of 


FIG.  85.  The  preparation  of  bromine  in 
the  laboratory 


the  water  in  the  flask  B,  which  is  partially  immersed  in 
ice  water.  On  heating,  the  bromine  distills  over,  and  is 
collected  in  the  cold  receiver.  The  equation  is 


2  NaBr+2  HSO+MnO 


HO+Br 


2.  Commercial  method.  Bromine  is  prepared  commercially 
from  the  waters  of  salt  wells  which  are  especially  rich  in 
bromides.  On  passing  a  current  of  electricity  through  such 
waters,  the  bromine  is  first  liberated.  Any  chlorine  liber- 
ated, however,  will  assist  in  the  reaction,  since  free  chlorine 
decomposes  bromides,  as  shown  in  the  equation 


2  NaBr  +  C1 


2  NaCl  +  Br 


When  the  water  containing  the   bromine  is  heated,  the 
liberated  bromine  distills  over  into  the  receiver. 


180  FIKST  COURSE  IN  CHEMISTRY 

Properties.  Bromine  is  a  dark-reddish-brown  liquid  about 
3  times  as  heavy  as  water ;  but  even  at  ordinary  temper- 
atures it  evaporates  rapidly,  forming  a  reddish-brown  gas 
very  similar  to  nitrogen  dioxide  in  appearance.  Its  vapor 
has  a  very  offensive  odor  and  is  most  irritating  to  the 
eyes  and  throat.  Bromine  is  somewhat  soluble  in  water, 
100  volumes  of  water  under  ordinary  conditions  dissolving 
1  volume  of  the  liquid.  It  is  readily  soluble  in  carbon 
disulfide,  forming  a  reddish  solution. 

Chemical  conduct  and  uses.  In  chemical  action  bromine 
is  very  similar  to  chlorine.  It  combines  directly  with  many 
of  the  same  elements  with  which  chlorine  unites,  but  with 
less  energy.  It  combines  with  hydrogen,  and  takes  away 
the  latter  element  from  some  of  its  compounds,  but  not  so 
readily  as  does  chlorine.  Its  bleaching  properties  are  also 
less  marked  than  those  of  chlorine. 

Bromine  finds  many  uses  in  the  manufacture  of  organic 
drugs  and  dyestuffs  and  in  the  preparation  of  bromides. 
Silver  bromide  is  extensively  used  in  photography,  and  the 
bromides  of  sodium  and  potassium  are  used  as  drugs. 

Hydrogen  bromide  (HBr).  When  sulfuric  acid  acts  upon 
a  bromide,  hydrogen  bromide  is  set  free : 

2  NaBr  +  H2SO4 >•  Na2SO4  +  2  HBr 

At  the  same  time,  some  bromine  is  liberated,  as  may  be  seen 
from  the  red  fumes  which  appear,  and  from  the  odor.  The 
explanation  of  this  is  found  in  the  fact  that  hydrogen 
bromide  is  much  less  stable  than  hydrogen  chloride  and  is 
therefore  more  easily  oxidized.  Concentrated  sulfuric  acid 
is  a  good  oxidizing  agent  (p.  154),  and  oxidizes  a  part  of 
the  hydrogen  bromide,  liberating  bromine  : 

H2S04  +  2  HBr ^2  H2O  4-  SO2  +  Br2 


THE  CHLOKIKE  FAMILY  181 

The  pure  compound  is  best  prepared  by  adding  water  to 
phosphorus  tribromide : 

PBr3  +  3  H20 *•  3  HBr  +  P(OH)3 

Properties.  Hydrogen  bromide  very  strikingly  resembles 
hydrogen  chloride  in  physical  and  chemical  properties.  The 
chief  point  in  which  it  differs  from  hydrogen  chloride  is 
in  the  fact  that  it  is  much  more  easily  oxidized,  so  that 
bromine  is  more  readily  set  free  from  it  than  chlorine 
is  from  hydrogen  chloride.  A  solution  of  hydrogen  bro- 
mide in  water  is  called  Jiydrobromic  acid.  The  salts  of 
hydrobromic  acid  are  known  as  bromides.  They  resemble 
the  chlorides  in  their  properties. 

IODINE 

Historical.  Iodine  was  discovered  in  1812  by  Courtois, 
in  the  ashes  of  certain  sea  plants.  Its  presence  was  revealed 
by  its  beautiful  violet  vapor. 

Occurrence.  Like  all  the  other  elements  of  the  chlorine 
family  it  is  not  found  in  the  free  state.  In  the  combined  state 
it  occurs  in  very  small  quantities  in  sea  water,  from  which 
it  is  absorbed  by  certain  sea  plants,  so  that  it  is  found  in 
their  ashes.  It  occurs  along  with  bromine  in  salt  deposits, 
and  is  also  found  in  Chile  as  a  constituent  of  the  enormous 
deposits  of  sodium  nitrate  (Chile  saltpeter). 

Preparation.  In  the  laboratory,  iodine  can  readily  be 
prepared  from  an  iodide  by  the  method  used  in  prepar- 
ing bromine,  sodium  iodide  being  substituted  for  sodium 
bromide. 

To  some  extent  iodine  is  prepared  commercially  by  burn- 
ing seaweed  (kelp)  at  a  low  temperature,  the  iodine  being 
left  in  the  ashes  as  sodium  iodide,  from  which  it  is  easily 


182  FIRST  COURSE  IN  CHEMISTRY 

obtained.  Chile  saltpeter  contains  a  very  small  percentage 
of  compounds  of  iodine,  and  most  of  the  iodine  of  commerce 
comes  from  the  liquors  obtained  in  the  purifying  of  this 
material. 

Properties.  Iodine  is  a  purplish-black,  shining,  heavy 
solid  which  crystallizes  in  brilliant  plates.  Even  at  ordinary 
temperatures  it  gives  off  a  beautiful  violet  vapor,  which 
increases  in  amount  as  heat  is  applied.  It  is  only  slightly 
soluble  in  water,  but  readily  dissolves  in  alcohol,  forming 
a  brown  solution  (tincture  of  iodine),  and  in  carbon  disul- 
fide,  forming  a  violet  solution.  The  element  has  a  strong, 
unpleasant  odor,  by  no  means  so  irritating  as  that  of  chlo- 
rine or  bromine. 

Chemical  conduct.  Chemically,  iodine  is  quite  similar  to 
chlorine  and  bromine,  but  is  still  less  active  than  bromine. 
Both  chlorine  and  bromine  displace  it  from  its  salts  : 


2KI  +  Br2  -  ^ 

2  KI  +  C12  -  >•  2  KC1  +  I2 

When  even  minute  traces  of  iodine  are  added  to  thin  starch 
paste  a  very  intense  blue  color  develops,  and  this  reaction 
forms  a  delicate  test  for  iodine.  Iodine  is  extensively  used 
in  medicine,  especially  in  the  form  of  a  tincture.  It  is 
also  largely  used  in  the  preparation  of  dyes  and  organic 
drugs.  lodoform,  a  substance  used  as  an  antiseptic,  has 
the  formula  CHI  . 

o 

Hydrogen  iodide  (HI).  This  gas  can  be  prepared  by  passing 
hydrogen  sulfide  into  water  in  which  iodine  is  suspended  : 

H2S  4-  12  —  *•  2  HI  +  S 

The  resulting  hydrogen  iodide  dissolves  in  the  water,  form- 
ing hydriodic  acid,  while  the  sulfur  is  precipitated. 


THE  CHLOEINE  FAMILY 


183 


Properties  and  uses.  Hydrogen  iodide  resembles  the  cor- 
responding compounds  of  chlorine  and  bromine,  but  is  even 
less  stable.  It  readily  decomposes  into  hydrogen  and  iodine, 
and  is  therefore  a  strong  reducing  agent. 

The  salts  of  hydriodic  acid,  the  iodides,  are,  in  general, 
similar  to  the  chlorides  and  bromides.  Potassium  iodide  is 
largely  used  in  medicine,  and  silver  iodide  in  photography. 


GAY-LUSSAC'S  .  LAW  OF  VOLUMES 

In  the  discussion  of  the  composition  of  hydrochloric  acid 
it  was  stated  that  1  volume  of  hydrogen  combines  with  1  vol- 
ume of  chlorine  to  form  2  volumes  of  hydrochloric  acid.  With 
bromine  and  iodine,  similar  combining  ratios  hold  good. 
These  facts  recall  the  simple  volume  relations  already  noted 
in  the  study  of  the  composition  of  steam  (p.  50)  and  am- 
monia (p.  127).  These  relations  may  be  represented  graphi- 
cally in  the  following  way,  the  squares  representing  equal 
volumes : 


In  the  early  part  of  the  past  century  the  distinguished 
French  chemist,  Gay-Lussac  (Fig.  23),  studied  the  volume 
relations  of  many  combining  gases,  and  concluded  that  simi- 
lar relations  always  hold.  His  observations  are  summed  up 
in  the  following  generalization,  known  as  the  law  of  volumes : 
When  two  gases  combine  chemically  there  is  ahvays  a  simple 
ratio  between  their  volumes,  and  also  between  the  volume  of 
either  one  of  them  and  that  of  the  product,  provided  it  is  a  gas. 
By  a  simple  ratio  is  meant,  of  course,  the  ratio  of  integer 
numbers,  as  1 :  2,  or  2  :  3. 


184  FIBST  COURSE  IN  CHEMISTRY 

EXERCISES 

1.  How  do  we  account  for  the  fact  that  liquid  hydrofluoric  acid 
is  not  an  electrolyte  ? 

2.  Why  is  hydrogen  fluoride  obtained  by  the  action  of  sulfuric 
acid  on  fluorite  ? 

3.  In  the  preparation  of  chlorine,  what  advantages  are  there  in 
treating  manganese  dioxide  with  a  mixture  of  sodium  chloride  and 
sulfuric  acid  rather  than  with  hydrochloric  acid? 

4.  Why  are  the  methods  of  preparation  used  in  the  laboratory 
likely  to  differ  from  those  used  commercially? 

5.  What  is  the  derivation  of  the  word  nascent? 

6.  What  substances  studied  are  used  as  bleaching  agents?    To 
what  is  the  bleaching  action  due  in  each  case  ? 

7.  What  does  the  term  muriatic  acid  signify  (see  dictionary)  ? 

8.  Upon  what  metals  would  you  expect  hydrochloric  acid  to  act  ? 
(See  Displacement  Series.) 

9.  What  is  the  meaning  of  the  phrase  aqua  regia? 

10.  A  solution  of  hydriodic  acid  turns  brown  on  standing.    Why? 

11.  From  their  behavior  toward  sulfuric  acid,  to  what  class  of 
agents  do  hydrobromic  and  hydriodic  acids  belong  ? 

12.  Give  the  derivation  of  the  names  of  the  elements  of  the 
chlorine  family. 

13.  What  is  formed  when  a  metal  dissolves  in  each  of  the  follow- 
ing :    nitric  acid ;    dilute  sulfuric  acid ;    concentrated  sulfuric  acid ; 
hydrochloric  acid  ;  aqua  regia  ? 

14.  In  what  respects  are  the  elements  included  in  the  chlorine 
family  similar? 

15.  What  weight  of  hydrogen   chloride  is  absorbed  by  1  liter 
of  water  under  standard  conditions?    Ans.  819.9  g.     What  weight 
of  sodium  chloride  is  necessary  to  prepare  this  weight  of  hydrogen 
chloride?    Ans.  1314.34  g. 

TOPICS  FOR  THEMES 

Scheele  (Thorpe,  Essays  in  Historical  Chemistry). 
Bleaching  (see  encyclopedia). 


CHAPTER  XX 
MOLECULAR  WEIGHTS ;   ATOMIC  WEIGHTS 

Introduction.  In  Chapter  IX  it  was  shown  that  from  the 
results  of  a  careful  analysis  of  a  compound  it  is  easy  to  cal- 
culate a  formula,  provided  we  adopt  the  simplest  one  possible. 
The  method  described  would  lead  to  the  formula  HO  for 
hydrogen  peroxide,  while  the  formula  we  have  adopted  is 
H2O2.  The  ratio  of  hydrogen  to  oxygen  is  the  same  in  these 
two  formulas,  and  to  decide  between  them  we  must  devise 
a  way  to  determine  the  weight  of  the  molecule  of  the  com- 
pound. If  this  weight  can  be  shown  to  be  17.008,  the 
correct  formula  is  HO  ;  if  it  is  34.016,  the  formula  is 
H2O2.  We  shall  now  turn  our  attention  to  the  problem  of 
determining  the  molecular  weights  of  compounds. 

Avogadro's  hypothesis.  At  the  close  of  the  last  chapter 
we  saw  that  two  gases  always  combine  in  some  simple  ratio 
by  volume,  and  that  the  volume  of  the  product  (if  it  is  a 
gas)  is  also  in  simple  ratio  to  that  of  the  other  two  volumes 
(law  of  volumes). 

These  relations  are  so  simple  and  so  unexpected  that  we 
at  once  feel  that  they  indicate  a  very  simple  ratio  between 
the  number  of  molecules  present  in  equal  volumes  of  gases. 
As  early  as  1811  the  Italian  physicist,  Avogadro,  suggested 
that  the  ratios  become  perfectly  intelligible  if  we  assume 
that  equal  volumes  of  any  two  gases  contain  the  same  number 
of  molecules.  This  generalization  is  known  as  the  hypothesis 
of  Avogadro,  and  it  is  in  complete  accord  with  all  we  have 
learned  about  gases  since  Avogadro's  time. 

185 


186  FIBST  COUKSE  IX  CHEMISTKY 

Avogadro's  hypothesis,  and  molecular  weights.  Assuming 
the  truth  of  Avogadro's  hypothesis,  we  have  a  simple  means 
of  deciding  upon  the  relative  weights  of  the  various  kinds 
of  molecules  ;  for  if  equal  volumes  of  two  gases  contain  the 
same  number  of  molecules,  the  weights  of  the  two  kinds  of 
molecules  must  be  in  the  same  ratio  as  the  weights  of  the 
two  volumes  made  up  of  these  molecules. 

For  example,  the  weight  of  a  liter  of  ammonia  is  0.7708  g. 
and  that  of  a  liter  of  hydrogen  chloride  is  1.6398  g.  These 
two  values  will  therefore  indicate  the  relative  weights  of 
the  two  kinds  of  molecules,  since  there  is  the  same  number 
of  each  in  a  liter.  If  we  adopt  some  one  gas  as  a  standard,  we 
can  readily  determine  the  weights  of  all  gaseous  molecules 
relatively  to  those  of  the  standard  gas.  Thus,  if  we  adopt 
ammonia  as  standard  (unity),  the  molecule  of  hydrogen 
chloride  is  2.14  times  as  heavy  as  the  standard. 

Oxygen  as  standard.  It  will  be  seen  that  the  gas  selected 
as  standard,  and  the  volume  chosen  for  comparison  will 
make  no  difference,  since  the  weights  are  all  relative  in  any 
case.  But  since  the  molecules  are  all  made  up  of  atoms,  it 
is  important  that  the  standard  chosen  for  atomic  weights 
should  be  in  accord  with  that  chosen  for  molecular  weights. 
For  many  reasons  oxygen  serves  best  for  atomic  weights,  and 
it  is  also  chosen  for  molecular  weights. 

The  oxygen  atom  taken  as  16.  In  Chapter  VI  we  saw 
that  oxygen  combines  with  hydrogen  in  the  ratio  1  :  7.94 
by  weight.  If  we  wish  to  take  oxygen  as  a  standard  for 
atomic  weights,  the  smallest  whole  number  we  can  assign 
to  it  while  keeping  hydrogen  at  least  unity  is  8,  hydrogen 
then  becoming  1.008.  Adopting  these  two  values,  the  for- 
mula of  water  will  be  HO.  There  is,  however,  good  reason 
for  thinking  that  there  are  two  atoms  of  hydrogen  in  a  mole- 
cule of  water.  For  example,  when  sodium  acts  upon  water, 


MOLECULAR  WEIGHTS;  ATOMIC  WEIGHTS    187 

exactly  one  half  of  the  hydrogen  is  given  off,  and  at  the 
same  time  sodium  hydroxide  is  formed,  which  contains  the 
other  half.  This  is  most  simply  expressed  in  the  equation 

H2O  +  Na *•  NaOH  +  H 

On  the  other  hand,  none  of  the  reactions  which  water 
undergoes  indicates  that  there  is  more  than  one  atom  of 
oxygen  in  a  molecule  of  water. 

If  there  are  two  atoms  of  hydrogen  and  one  of  oxygen  in 
water,  then  to  keep  the  hydrogen  atom  at  least  unity,  we 
must  double  the  value  of  oxygen,  making  it  16  instead  of 
8.  The  formula  of  water  then  becomes  H2O. 

The  oxygen  molecule  taken  as  32.  Since  we  wish  to  use 
oxygen  gas  as  a  standard  with  which  to  compare  other 
gases,  it  is  important  to  assign  a  weight  to  the  oxygen 
molecule  that  will  keep  the  atom  equal  to  16.  Now  when 
hydrogen  and  oxygen  combine  to  form  steam,  we  have  the 
equation  (p.  50)  : 

2  vol.  hydrogen  + 1  vol.  oxygen >•  2  vol.  steam 

Let  us  suppose  that  the  1  volume  of  oxygen  contains 
100  molecules ;  then  the  2  volumes  of  steam  must  con- 
tain 200  molecules  (Avogadro's  hypothesis).  But  each  of 
these  200  molecules  must  contain  at  least  one  atom  of  oxy- 
gen, or  200  in  all,  and  these  200  atoms  came  from  100 
molecules  of  oxygen.  Therefore  each  molecule  of  oxygen 
must  contain  at  least  two  atoms  of  oxygen. 

Evidently  this  reasoning  merely  shows  that  there  are  at 
least  two  atoms  in  the  oxygen  molecule.  There  may  be  more 
than  that,  but  as  there  is  no  evidence  that  this  is  so,  we 
assume  that  each  oxygen  molecule  contains  only  two  atoms. 
If,  then,  we  wish  to  retain  the  value  16  for  the  oxygen  atom, 
we  must  adopt  32  for  the  value  of  the  oxygen  molecule. 


188  FIKST  COUESE  IN  CHEMISTRY 

Molecular  weights  from  weight  of  i  liter.  We  have  now 
devised  a  method  of  determining  how  much  heavier  one 
kind  of  a  molecule  is  than  another,  and  have  fixed  upon 
the  weight  of  one  standard  molecule  (oxygen),  with  which 
all  others  can  be  compared.  The  determination  of  molec- 
ular weights  now  becomes  easy.  For  example,  1  liter  of 
oxygen  weighs  1.429  g.,  while  1  liter  of  hydrogen  chloride 
weighs  1.6398  g.  The  ratio  between  the  weights  of  the 
two  kinds  of  molecules  is  therefore  1.429  : 1.6398.  To  com- 
pare the  hydrogen  chloride  molecule  with  oxygen  taken  as 
32,  we  need  only  solve  the  proportion:  1.429  : 1.G398  : :  32  :  x. 
The  molecular  weight  of  hydrogen  chloride  (V)  is  there- 
fore 36.7. 

Gram-molecular  volume  equals  22.4  liters.  Having 
adopted  32  as  the  standard  for  oxygen,  it  is  of  interest  to  find 
the  volume  occupied  by  the  gram-molecular  weight  of  this 
gas,  namely,  32  g.  A  simple  computation  from  the  weight 
of  1  liter  shows  this  volume  to  be  22.4  liters.  If  we  con- 
struct a  vessel  of  exactly  this  content  and  fill  it  with  oxy- 
gen gas,  it  will  contain  just  enough  molecules  of  oxygen 
to  weigh  32  g.,  which  is  our  standard  weight  for  oxygen. 

If,  now,  we  replace  the  oxygen  by  another  gas,  say,  hydro- 
gen chloride,  we  shall  have  the  same  number  of  molecules 
present.  The  weight  of  hydrogen  chloride  filling  the  vessel 
is  36.45  g.  But  since  there  are  the  same  number  of  mole- 
cules, the  values  32  and  36.45  must  represent  the  relative 
weights  of  the  two  kinds  of  molecules.  In  like  manner,  the 
weight  of  22.4  liters  of  any  gas  will  give  a  number  which 
expresses  the  weight  of  a  molecule  of  that  gas  compared 
with  the  molecule  of  oxygen  taken  as  the  standard.  These 
relations  are  illustrated  in  Fig.  86.  We  therefore  reach  the 
following  simple  rule:  The  molecular  weight  of  any  gas  may 
be  found  by  determining  the  weight  in  grams  of  22.4  liters 


MOLECULAR  WEIGHTS;  ATOMIC  WEIGHTS    189 

of  the  gas.  The  volume  22.4  liters  is  called  the  gram- 
molecular  volume  of  gases.  Owing  to  the  fact  that  most 
gases  do  not  exactly  conform  to  any  of  the  gas  laws,  the 
weight  of  22.4  liters  of  a  gas  is  not  its  precise  molecular 
weight,  but  is  very  close  to  it. 

It  is  evident  that  this  method  only  applies  to  elements  or 
compounds  that  are  gases  or  can  be  converted  into  the  gas- 
eous state.  For  all  others  different  methods  are  known,  but 
their  discussion  is  beyond  the  scope  of  an  elementary  text. 


</ 

/ 

/ 

^          ^ 

^          ^ 

/ 

^ 

/ 

22.4 
LITERS 

22.4 
LITERS 

/ 

22.4 
LITERS 

/ 

22.4 
LITERS 

O232g.  HC1  36.45  g.  NH3  17.034  g.  H2O18.01Gg. 

FIG.  86.  The  weight  of  22.4  liters  of  various  gases 

Molecular  weights  of  the  elements.  When  we  determine 
the  weight  of  22.4  liters  of  the  various  elementary  gases, 
we  reach  some  interesting  conclusions.  Experiment  shows 
that  the  molecular  weight  of  many  of  them,  such  as  nitro- 
gen, hydrogen,  chlorine,  and  bromine  give  values  which  are 
twice  the  atomic  weights,  so  that  in  these  cases  the  molecule 
contains  two  atoms  (p.  69).  In  the  case  of  the  metals,  so 
far  as  their  vapors  have  been  studied,  the  molecular  weight 
and  the  atomic  weight  are  the  same,  so  that  the  molecular 
consists  of  a  single  atom.  The  molecule  of  ozone  contains 
three  atoms  of  oxygen,  so  that  its  formula  is  O3 ;  while  the 
molecules  of  phosphorus  and  arsenic  contain  four  atoms, 
giving  the  formulas  P4  and  As4. 

Selection  of  atomic  weights  from  combining  weights.  It 
is  now  easy  to  determine  which  multiple  of  the  combining 
weight  shall  be  adopted  as  the  correct  atomic  weight  —  a 
problem  which  we  left  unsolved  in  Chapter  VIII.  (The 


190 


FIRST  COUKSE  IN  CHEMISTRY 


student  should  at  this  point  carefully  review  page  70.) 
The  mode  of  procedure  will  be  understood  most  readily  by 
an  example  ;  so  let  us  suppose  that  we  have  found  the  com- 
bining weight  of  nitrogen  to  be  7.005  and  that  we  wish  to 
decide  whether  this  value  or  some  simple  multiple,  14.01 
or  21.015,  is  the  atomic  weight. 

We  first  determine  the  weight  of  22.4  liters  of  a  number 
of  gaseous  compounds  which  we  know  to  contain  nitrogen. 
These  values  are  given  in  the  first  column  of  the  table. 


NAME  OF  GASEOUS 
COMPOUND 

MOLECULAR 
WEIGHT 

(22.4  LITERS) 

PERCENTAGE  OF 
NITROGEN  BY 
EXPERIMENT 

PART  OF  MOLECU- 
LAR WEIGHT  DUE 
TO  NITROGEN 

Nitrogen  gas  . 
Nitrous  oxide  . 

27.95 

41.1:} 

100.00 
63.70 

27.95 
28.11 

Nitric  oxide     .     .     . 

10.00 

46.74 

14.02 

Ammonia    .... 

17.05 

82.28 

14.03 

Nitric  acid 

Q3.75 

22.27 

14.30 

We  next  make  a  careful  analysis  of  each  of  these  com- 
pounds to  ascertain  the  percentage  of  nitrogen  present, 
placing  the  values  obtained  in  the  second  column.  If  we 
multiply  the  molecular  weight  of  each  compound  by  the 
percentage  of  nitrogen,  the  product  will  be  the  portion  of 
the  molecular  weight  due  to  nitrogen.  But  since  the  mole- 
cules are  made  up  of  atoms,  the  part  of  a  molecule  due  to 
nitrogen  must  represent  the  mim  of  the  weights  of  the  nitro- 
gen atoms  present.  We  notice  that  the  numbers  in  the  last 
column  are  either  very  near  to  14  or  to  twice  14  and  that 
none  are  near  7.  If  we  examine  a  large  number  of  nitrogen 
compounds,  it  is  reasonable  to  expect  that  we  should  find 
some  containing  only  one  atom,  and  since  we  find  none 
which  give  a  value  of  less  than  14,  we  assume  that  this  and 
not  7  or  21  or  28  represents  the  weight  of  a  nitrogen  atom. 


MOLECULAR  WEIGHTS;  ATOMIC  WEIGHTS    191 

Accurate  determination  of  atomic  weights.  The  weight 
of  a  given  volume  of  a  gas  is  difficult  to  determine  with 
great  precision,  and  in  consequence  the  molecular  weights 
of  gases  as  determined  by  experiment  are  usually  subject 
to  a  very  considerable  error.  The  portion  of  nitrogen  in 
22.4  liters  of  the  various  gases  is  therefore  a  little  uncertain, 
as  will  be  seen  from  the  values  in  the  last  column  (above). 
All  these  figures  tell  us  is  that  the  true  value  is  very 
near  14.  The  combining  iveiyht  can  be  very  accurately  deter- 
mined by  the  analysis  of  any  of  these  compounds,  and  is 
found  to  be  7.005.  It  is  therefore  evident  that  the  accurate 
atomic  weight  is  twice  this  value,  namely,  14.01. 

Summary.  These,  then,  are  the  steps  which  must  be 
taken  to  establish  the  atomic  weight  of  an  element. 

1.  Determine  the  combining  weight  accurately  by  analysis. 

2.  Determine  the  weight  of  22.4  liters  of  a  large  number  of 
gaseous  compounds  of  the  element,  and,  by  analysis,  the  part 
of  the  molecular  weights  clue  to  the  element.   The  smallest 
number  so  obtained  will  be  the  approximate  atomic  weight. 

3.  Multiply  the  combining  weight_byjthe  integer  (1,  2, 
or  3)  which  will  give  a  number  close  to  the  approximate 
atomic  weight.   The  number  so  obtained  will  be  the  precise 
atomic  weight. 

Equations  and  volumes  of  gases.  If  we  have  an  equation 
which  expresses  a  reaction  in  which  gaseous  molecules  take 
part,  we  may  make  use  of  Avogadro's  hypothesis  to  predict 
the  volume  changes  which  will  accompany  the  reaction. 
For  example,  the  equation 


states  that  1  gram-molecular  weight  of  hydrogen  combines 
with  1  gram-molecular  weight  of  chlorine  to  give  2  gram- 
molecular  weights  of  hydrogen  chloride.  Now  all  of  these 


192  FIEST  COUESE  IN  CHEMISTRY 

substances  are  gases,  and  a  gram-molecular  weight  of  every 
gas  occupies  the  same  volume,  namely,  22.4  liters.  Conse- 
quently, 1  volume  of  hydrogen  will  combine  with  1  volume 
of  chlorine  to  give  2  volumes  of  hydrogen  chloride,  and 
there  will  be  no  change  in  the  volume  due  to  the  reaction 
(save  as  occasioned  by  the  heat  given  off).  The  coefficients 
of  the  molecules  therefore  indicate  the  proportion  by  volume  in 
which  gases  take  part  in  reactions. 

Weight  of  a  liter  of  a  gas.  We  have  found  that  a  gram- 
molecular  weight  of  any  gas  occupies  22.4  liters.  If  we 
know  the  molecular  weight  of  a  gas,  we  can  at  once  deduce 
the  weight  of  a  liter  of  the  gas.  For  example,  the  molec- 
ular weight  of  acetylene  (C  H  )  is  26.016.  This  means 
that  26.016  g.  occupy  22.4  liters.  Consequently  1  liter  will 
weigh  26.016  -=-22.4  =  1.1614  g.  In  general,  to  find  the 
weight  of  a  liter  of  any  gas,  divide  its  molecular  weight 
by  22.4.  The  value  so  obtained  will  be  close  enough  to  the 
experimental  value  for  all  practical  purposes. 

EXERCISES 

1.  What  are  the  relative  weights  of  the  molecules  of  hydrogen 
and  hydrogen  chloride,  as  deduced  from  a  weight  of  1  liter  of  each 
of  these  gases?  Ans.    As  1  is  to  18.34. 

2.  Natural  gas  is  largely  composed  of  marsh  gas  (CII4).    When 
this  burns,  the  equation  for  the  reaction  is  as  follows : 

CII4  +  2  O2 >•  CO2  +  2  II2O 

In  burning  100  cu.  ft.  of  this  gas,  what  volume  of  oxygen  is  consumed? 
Ans.  200  cu.  ft.  What  is  the  volume  of  the  carbon  dioxide  formed? 
Ans.  100  cu.  ft. 

3.  Why  write  2  O2  rather  than  4  O  in  problem  2  ? 

4.  The  molecular  weight  of  nitric  oxide  is  30.01.    What  is  the 
approximate  weight  of  1  liter  of  the  gas?    Ans.  1.34  g. 

5.  A  compound  was  found  to  have  the  composition  II  =  5.91%, 
S  =  94.08%;  mol.  wt.  =  34.1.    Calculate  its  formula.    Ans.  II2S. 


CHAPTER  XXI 


CARBON  AND   SOME   OF  ITS   SIMPLER  COMPOUNDS 

Occurrence.  In  the  uncombined  state  carbon  is  found 
in  nature  in  several  forms.  The  diamond  is  practically 
pure  carbon,  while  graphite  and  the  various  forms  of  coal 
all  contain  more  or  less  free  carbon.  The  element  also 
occurs  abundantly  in  the 
form  of  compounds.  Car- 
bon dioxide  is  its  most  fa- 
miliar gaseous  compound. 
Natural  gas  and  petroleum 
are  largely  compounds  of 
carbon  and  hydrogen.  The 
carbonates,  especially  cal- 
cium carbonate  (limestone), 
constitute  great  strata  of 
rocks,  and  are  found  in 
almost  every  locality.  All 
living  organisms,  both  plant  and  animal  (p.  89),  contain 
a  large  percentage  of  this  element,  and  the  number  of  its 
compounds  which  go  to  make  up  the  vast  variety  of  ani- 
mate nature  is  almost  limitless.  In  the  free  state  carbon 
occurs  in  both  the  crystalline  and  the  amorphous  form. 

Crystalline   carbon.    Crystalline   carbon    occurs  in  two 
forms  —  the  diamond  and  graphite. 

1.  Diamond.    Diamonds  are  found  in  certain  localities  in 
South  Africa,  the  East  Indies,  and  Brazil.    The  crystals  as 

193 


FIG.  87.   The  Cullinan  diamond  in 
its  original  condition  (one  half  nat- 
ural size) 


194 


FIRST  COURSE  IN  CHEMISTRY 


FIG.  88.  The  Kohinoor  diamond 
after  being  cut  (natural  size) 


found  are  usually  covered  with  a  rough  coating.  These  are 
cut  so  as  to  bring  out  the  brilliancy  of  the  gem.  Diamond 

cutting  is  carried  on  most  exten- 
sively in  Holland. 

Fig.  87  is  a  photograph  (one  half 
natural  size)  of  the  largest  diamond 
ever  found,  in  its  original  condi- 
tion. It  is  known  as  the  Cullinan 
diamond,  and  was  presented  to 
King  Edward  VII  by  the  Trans- 
vaal government.  Fig.  88  is  a 
photograph  (natural  size)  of  an- 
other very  famous  diamond  (the 
Kohinoor),  in  finished  form. 

The  density  of  the  diamond  is  8.5,  and,  though  brittle,  it 
is  one  of  the  hardest  of  substances.  Few  chemical  reagents 
have  any  action  on  it,  but  when  heated  in  oxygen  or  the 
air,  it  blackens  and  burns,  form- 
ing carbon  dioxide. 

Artificial  preparation  of  diamonds. 

Many  attempts  have  been  made  to 
produce  diamonds  artificially.  For 
a  long  time  these  ended  in  failure, 
graphite  and  not  diamonds  being 
the  product  obtained,  but  in  1893 
the  French  chemist,  Moissan  (Fig. 
77),  in  his  study  of  chemistry  at  high 
temperatures,  finally  succeeded  in 
making  some  small  ones.  He  ac- 
complished this  by  dissolving  car- 
bon in  melted  iron  and  plunging 

the  crucible  containing  the  mixture  into  water,  as  shown  in 
Fig.  89.  Under  these  conditions  the  carbon  crystallized  in  the 
iron  in  the  form  of  the  diamond.  The  diamonds  were  then  freed 
from  the  metal  by  dissolving  away  the  iron  in  hydrochloric  acid.. 


Fi<;.  80.    The  artificial  prep- 
aration of  the  diamond 


CAKBON  AND  ITS  SIMPLER  COMPOUNDS    195 

2.  Graphite.  This  form  of  carbon  is  found  in  large  quanti- 
ties, especially  in  Ceylon,  Siberia,  and  in  some  parts  of  the 
United  States  and  Canada.  Large  quantities  are  also  made 
commercially  by  heating  hard  coal  to  a  high  temperature.  It 
is  a  glistening  black  substance,  very  soft,  and  greasy  to  the 
touch.  Its  density  is  about  2.15.  It  is  used  in  the  manufac- 
ture of  lead  pencils  and  crucibles,  as  a  lubricant,  and,  in  the 
form  of  a  polish  or  a  paint,  as  a  protective  covering  for  iron. 

Commercial  production  of  graphite.  The  process  consists  in  heat- 
ing hard  coal  in  large  electric  furnaces  about  40  ft.  in  length, 
a  longitudinal  section  of  one  of  which  is  shown  in  Fig.  90. 


FIG.  90.    Electric  furnace  for  the  production  of  graphite 

The  electrodes  A  are  made  of  graphite.  The  furnace  is  nearly 
filled  with  the  coarse  grains  of  coal  B.  Since  the  coal  is  a 
poor  conductor,  there  is  placed  in  the  center  of  the  charge  a 
core  C  of  carbon,  which  serves  to  conduct  the  current  through 
the  charge.  The  charge  is  covered  with  a  mixture,  D,  of  sand 
and  carbon  (or  similar  materials),  which  serves  to  exclude  the 
air.  An  alternating  current  is  supplied  by  the  generator  G. 
Under  the  influence  of  the  intense  heat  produced  by  the  current, 
the  carbon  is  changed  into  graphite.  Prepared  in  this  way,  the 
product  is  uniform  in  composition  and  free  from  grit,  and  is 
therefore  superior  to  the  natural  product. 

Amorphous  carbon.  Pure  amorphous  carbon  is  best  pre- 
pared by  heating  sugar  (C12H22On)  m  the  absence  of  air. 
The  hydrogen  and  oxygen  present  are  expelled,  largely  in 


196  FIRST  COURSE  IN  CHEMISTRY 

the  form  of  water,  and  pure  carbon  remains.  Among  the 
numerous  substances  that  contain  amorphous  carbon,  the 
following  may  be  mentioned : 

1.  Coal  and  coke.    The  various  forms  of  coal  were  formed 
from  vast  accumulations  of  vegetable  matter.    In  hard  coal, 
or  anthracite,  nearly  all  the  carbon  present  is  in  the  uncom- 
bined  state ;  while  in  soft,  or  bituminous  coal,  a  considerable 
portion  of  the  carbon  present  is  combined  with  hydrogen, 
oxygen,  nitrogen,  and  sulfur.    When  soft  coal  is  heated  in 
the  absence  of  air  (see  coal  gas,  p.  210),  complex  changes 
occur,  resulting  in  the  formation  of  various  useful  com- 
pounds of  carbon,  which  are  given  off  in  the  form  of  gases  and 
vapors,  while  the  mineral  matter  and  free  carbon  remain  and 
constitute  ordinary  coke.    The  matter  which  escapes  when 
coal  is  heated  in  the  absence  of  air  is  known  as  volatile  mat- 
ter.   In  hard  coal  the  volatile  matter  averages  from  5  per 
cent  to  8  per  cent,  while  in  soft  coal  it  averages  from  30 
per  cent  to  35  per  cent.    When  coal  is  burned,  the  mineral 
matter  present  is  left  in  the  form  of  ash. 

2.  Charcoal.    This  is  prepared  from  wood  just  as  coke  is 
prepared  from  coal.    The  volatile  matter  expelled  contains 
many  valuable  substances,  such  as  wood  alcohol  and  acetic 
acid,  which  are  obtained  commercially  in  this  way.  Formerly 
much  of  this  volatile  matter  was  allowed  to  escape,  but  at 
present  an  increasing  amount  of  charcoal  is  prepared  in 
such  a  way  that  the  volatile  matter  is  condensed  and  saved, 
as  in  the  heating  of  coal.    Both  charcoal  and  coke  are  used 
as  fuels,  and  they  are  especially  useful  in  reducing  metals 
from  their  oxides,  as  will  be  noted  later. 

Modern  methods  for  the  production  of  charcoal.  Iron  cars  are 
loaded  with  wood  A,  A  (Fig.  91)  and  run  into  the  retort  B.  The 
retort  is  then'made  air-tight  and  heated  slowly  for  twenty-four 
hours  by  the  fires  F,  F.  The  volatile  products  escape  through 


CARBON  AND  ITS  SIMPLER  COMPOUNDS    197 

the  pipes  C,  C  and  then  pass  into  the  condensers  D,  D.  Here 
those  portions  which  are  liquid  at  ordinary  temperatures,  such 
as  wood  alcohol  and  acetic  acid,  are  condensed,  while  the  gase- 
ous products  are  led  back  into  the  furnace  and  burned.  When 
all  the  volatile  matter  has  been  expelled  in  this  way,  the  cars 
containing  the  charcoal  are  run  into  cooling  chambers,  and  their 
place  in  the  retort  is  taken  by  other  cars  loaded  with  wood. 


FIG.  91.  A  modern  plant  for  the  production  of  charcoal 

3.  Bone  black,  or  animal  charcoal.   This  is  made  by  charring 
bones  and  animal  refuse.    It  consists  of  very  finely  divided 
carbon  and  of  calcium  phosphate,  and  is  especially  useful 
for  removing  coloring  matter  in  the  refining  of  sugar. 

4.  Lampblack.   Lampblack  is  a  product  of  the  imperfect 
combustion  of  carbonaceous  fuels,  such  as  oil  and  gas.    It 
is  used  in  making  indelible  inks,  printer's  ink,  and  black 
varnishes. 

Properties.  While  the  various  forms  of  carbon  differ  in 
many  properties,  especially  in  hardness,  yet  they  are  all  odor- 
less, tasteless  solids,  insoluble  in  water,  and  characterized  by 


198  FIRST  COURSE  IN  CHEMISTRY 

their  stability  towards  heat.  Only  in  the  intense  heat  of  the 
electric  arc  does  carbon  volatilize  to  any  considerable  extent. 

Chemical  conduct.  At  ordinary  temperatures  carbon  is  a 
very  inert  substance,  but  at  a  higher  temperature  it  com- 
bines directly  with  most  of  the  elements.  Because  of  its 
strong  affinity  for  oxygen  it  is  an  excellent  reducing  agent. 
Its  compounds  with  the  metals  are  called  carbides.  One  of 
the  most  important  of  these  is  calcium  carbide  (CaC2),  which 
is  used  in  the  preparation  of  acetylene.  When  carbon  or 
a  substance  containing  it,  such  as  wood  or  coal,  burns,  the 
element  combines  with  oxygen  to  form  either  carbon  dioxide 
(CO2)  or  carbon  monoxide  (CO).  Both  of  these  oxides 
are  colorless  gases. 

Carbon  dioxide  ;  preparation.  Attention  has  already  been 
called  (p.  89)  to  the  presence  of  this  gas  in  the  atmosphere 
and  to  the  various  natural  processes  by  means  of  which  it 
is  formed.  In  the  laboratory  it  is  always  prepared  by  the 
action  of  an  acid  upon  a  carbonate  (usually  calcium  car- 
bonate, in  the  form  of  marble).  This  reaction  might  be 
expected  to  produce  carbonic  acid,  thus: 

CaC08  +  2  HC1 >•  CaCl2  +  H2CO8 

Carbonic  acid  is  very  unstable,  however,  and  decomposes 
into  its  anhydride,  CO2,  and  water,  thus: 

H2CO3 ^H20  +  CO2 

In  the  preparation  of  carbon  dioxide,  pieces  of  marble  are 
placed  in  the  flask  A  (Fig.  70).  Hydrochloric  acid  is  added 
drop  by  drop  through  the  funnel  tube  B.  The  carbon  dioxide 
escapes  through  C  and,  being  heavier  than  air,  collects  in  the 
cylinder,  as  shown  in  the  figure. 

Properties.  Carbon  dioxide  is  a  colorless,  almost  odorless 
gas,  whose  density  is  1.5.  Its  weight  may  be  inferred  from 
the  fact  that  it  can  be  poured  like  water  from  one  vessel 


CAKBON  AND  ITS  SIMPLER  COMPOUNDS    199 


downward  into  another.  At  15°  and  under  ordinary  pres- 
sure, 1  volume  of  water  dissolves  1  volume  of  the  gas.  It 
is  rather  easily  condensed  to  a  colorless  liquid,  which  is 
slightly  lighter  than  water  and  boils  at  —  78.2°. 

Liquid  and  solid  carbon  dioxide.  The  commercial  carbon  diox- 
ide compressed  in  steel  cylinders  is  under  such  great  pressure 
that  it  is  largely  in  the  liquid  state.  When  the  pressure  is  re- 
moved, the  rapid  vaporization  of  the  liquid  reduces  the  tem- 
perature sufficiently  to  freeze  a  portion  of  the  escaping  liquid 
to  a  snowlike  solid  (Fig.  92).  Cylinders  of  liquid  carbon  dioxide 
are  inexpensive,  and  should 
be  available  in  every  school. 
The  commercial  supply  of  this 
gas  is  obtained  largely  from 
fermentation  processes,  espe- 
cially from  breweries. 

To  prepare  the  solid  carbon 
dioxide,  the  cylinder  should 
be  placed  across  the  table  and 
supported  in  such  a  way  that 
the  stopcock  end  is  several 
inches  lower  than  the  other 
end.  A  loose  bag  is  made  by 
holding  the  corners  of  a  piece 

of  cloth  around  the  neck  of  the  stopcock.  The  stopcock  is  then 
turned  on  so  that  the  liquid  rushes  out  in  large  quantities.  Very 
quickly  a  considerable  quantity  of  the  snow  collects  in  the  cloth. 

Chemical  conduct.  Carbon  dioxide  is  neither  combustible 
nor  a  supporter  of  combustion.  When  passed  over  carbon 
heated  to  a  high  temperature,  carbon  monoxide  is  formed  : 

CO2  +  C +  2  CO 

Uses.  The  relation  of  carbon  dioxide  to  plant  life  has 
been  discussed  in  a  previous  chapter  (p.  89).  Water  highly 
charged  with  carbon  dioxide  is  used  for  making  soda  water 


FIG.  92.   Carbon  dioxide  in  the 
solid  form 


200 


FIRST  COURSE  IN  CHEMISTRY 


and  similar  beverages.  Ordinary  soda  water  consists  of  dif- 
ferent flavoring  extracts,  to  which  is  added  water  highly 
charged  with  carbon  dioxide.  The  pressure  being  removed, 
the  excess  of  gas  escapes,  producing  effervescence.  Carbon 
dioxide  is  also  used  as  a  fire  extinguisher.  Most  of  the 
portable  fire  extinguishers  are  simply  devices  for  generating 
large  volumes  of  the  gas.  A  comparatively  small  percentage 
of  the  gas  in  air  is  sufficient  to  smother  a  flame. 

A  familiar  type  of  portable  fire  ex- 
tinguisher is  shown  in  Fig.  93.  The 
liquid  A  is  a  solution  of  sodium  car- 
bonate in  water.  The  bottle  B  contains 
sulfuric  acid.  In  case  of  fire  the  appa- 
ratus is  grasped  by  the  handle  C,  and 
the  knob  D  is  pushed  in  by  tapping  it 
against  the  floor.  This  breaks  the  bottle 
containing  the  sulfuric  acid,  which  at 
once  reacts  with  the  sodium  carbonate, 
generating  carbon  dioxide.  The  pressure 
of  the  gas  forces  the  water  out  through 
the  nozzle  E.  While  the  volume  of  water 
so  obtained  is  not  large,  it  is  very  effec- 
tive as  a  fire  extinguisher,  because  of 
the  carbon  dioxide  accompanying  it. 


c 
FIG.  93. 


A  portable  fire 
extinguisher 


Carbonic  acid  (H2C03).  This  acid  is  unstable  and  is  known 
only  in  the  form  of  a  very  dilute  solution.  This  solution 
is  most  readily  prepared  by  passing  carbon  dioxide  into 
water : 


HO  +  CO. 


HOO. 


(1) 


The  resulting  solution  neutralizes  bases  and,  in  general, 
possesses  the  properties  of  a  weak  acid. 

The  volume  of  carbon  dioxide  absorbed  in  pure  water  is 
relatively  small.  If,  however,  the  water  contains  a  base  (such 
as  sodium  hydroxide)  in  solution,  the  carbonic  acid  formed 


CAKBON  AND  ITS  SIMPLER  COMPOUNDS    201 

according  to  equation  (1)  reacts  with  the  base  to  form  the 
corresponding  carbonate : 

H2C03  +  2  NaOH  ^  Na2C03  +  2  HaO  (2) 

The  removal  of  the  carbonic  acid  results  in  the  union  of  more 
carbon  dioxide  and  water,  according  to  equation  (1),  so  that  the 
absorption  of  carbon  dioxide  will  continue  until  the  base  has 
been  changed  into  the  corresponding  carbonate. 

Salts  of  carbonic  acid  ;  carbonates.  The  carbonates  form 
an  important  class  of  salts.  Limestone,  shells,  and  marble 
are  largely  calcium  carbonate  (CaCO3),  common  washing 
soda  is  sodium  carbonate  (NagCO8),  while  baking  soda  is 
the  acid  carbonate  (NaHCO3).  The  carbonates  are  readily 
acted  upon  by  acids  liberating  carbon  dioxide  (see  prepa- 
ration of  carbon  dioxide).  When  heated  to  a  high  temper- 
ature, most  of  the  carbonates  readily  decompose,  forming 
carbon  dioxide  and  an  oxide  of  the  metal.  Thus,  lime 
(calcium  oxide,  CaO)  is  made  by  strongly  heating  calcium 
carbonate:  CaCO. ->  CaO  +  CO, 

Action  of  carbon  dioxide  on  calcium  hydroxide.  If  carbon 
dioxide  is  passed  into  a  solution  of  calcium,  hydroxide 
(ordinary  lime  water),  calcium  carbonate  is  formed  and, 
being  insoluble,  precipitates : 

H20  +  C00  — >•  H2C03 
Ca(OH)2  +  H2C03 >-  CaC03  +  2  H2O 

Advantage  is  taken  of  this  reaction  in  testing  for  the  pres- 
ence of  carbon  dioxide.  For  example,  if  the  air  exhaled 
from  the  lungs  is  blown  through  clear  limewater,  the  solu- 
tion soon  becomes  milky,  owing  to  the  precipitation  of  the 
white  solid,  calcium  carbonate  thus  proving  the  presence  of 
carbon  dioxide  in  the  exhaled  air. 


202  FIRST  COURSE  IN  CHEMISTRY 

Carbon  monoxide.  This  gas  is  formed  whenever  carbon 
is  burned  in  a  limited  supply  of  oxygen: 

2C-fO2 — ^2  CO 

It  is  often  formed  in  stoves  when  the  air  draft  is  shut  off, 
especially  when  hard  coal  is  used  as  a  fuel.  Since  the  gas  is 
very  poisonous,  care  should  be  taken  that  the  pipes  and  chim- 
neys are  not  closed  in  any  way ;  otherwise  the  gas  may  escape 
into  the  room  and  cause  the  death  of  those  inhaling  it. 

In  the  laboratory  the  gas  is  usually  prepared  by  heating 
formic  acid  (H2CO2)  or  oxalic  acid  (H2C2O4)  with  sulfuric 
acid,  which  is  added  to  absorb  the  water  formed : 

H2C02— >H20  +  CO 

H2C2°4  ^  H2°   +  C°2  +  C° 

Properties  and  chemical  conduct.  Carbon  monoxide  is  a 
colorless  gas  slightly  lighter  than  air.  It  is  a  very  treacher- 
ous poison,  since  it  is  almost  odorless. 

It  is  interesting  to  note  that  birds  are  very  sensitive  to  this 
gas.  In  mine  explosions  carbon  monoxide  is  always  formed, 
and  rescuers  often  carry  canaries  with  them,  the  death  of  the 
birds  warning  the  rescuers  of  their  own  peril. 

Chemically  carbon  monoxide  is  quite  active.  It  combines 
readily  with  oxygen,  and  burns  in  air  with  a  characteristic 
pale-blue  flame  (often  observed  in  the  combustion  of  hard 
coal),  forming  carbon  dioxide 

2  CO  +  02  — >-  2  C02 

It  reduces  metallic  oxides  such  as  copper  oxide   (CuO), 
forming  the  metal  and  carbon  dioxide : 

CO  +  CuO *•  CO2  +  Cu 

Because  of  this  property  carbon  monoxide  is  often  used  as 
a  reducing  agent  in  liberating  the  metals  from  their  oxides. 


CARBON  AND  ITS  SIMPLER  COMPOUNDS    203 

Cyanogen  (C2N2)  and  hydrocyanic  acid  (HNC).  At  high 
temperatures  carbon  unites  with  nitrogen  to  form  the  color- 
less, very  poisonous  gas,  cyanogen  (C2N2).  With  hydrogen 
and  nitrogen  it  forms  hydrocyanic  acid,  often  called  prus- 
sic  acid.  This  is  a  colorless  liquid  boiling  at  26°  and  is 
one  of  the  most  poisonous  substances  known.  Its  vapor 
is  often  used  to  kill  insects.  The  salts  of  prussic  acid 
are  known  as  cyanides.  They  are  likewise  very  poisonous. 
Sodium  cyanide  (NaNC)  and  potassium  cyanide  (KNC) 
are  white  solids.  Their  solutions  readily  dissolve  gold, 
and  are  used  in  extracting  gold  from  its  ores. 

The  hydrocarbons.  Carbon  and  hydrogen  unite  to  form 
a  great  many  compounds.  These  are  known  as  the  hydro- 
carbons. Their  importance  may  be  inferred  from  the  fact 
that,  mixed  together  in  varying  proportions,  they  constitute 
such  valuable  substances  as  natural  gas,  gasoline,  kerosene, 
vaseline,  and  paraffin. 

In  order  to  simplify  the  study  of  the  hydrocarbons,  it 
is  convenient  to  arrange  them  in  groups,  or  series.  The 
most  important  of  these  is  the  methane  series.  In  the 
following  table  are  given  the  names,  formulas,  and  boil- 
ing points  of  some  of  the  members  of  this  series  having  a 
small  number  of  carbon  atoms : 

BOILING  BOILING 

POINT  POINT 

Methane  (CII4)  .  .  -  160°  Pentane  (C5II12)      .     .     +  36° 

Ethane  (C2IIG)  .  .  -  93°  Hexane  (Ccllj)      .     .      +69° 

Propane  (C3II8)  .  .  -  45°  Heptane  (C7II1C)     .     .     +98° 

Butane  (C4H10)  .  .  +1°  General  formula  (CnH2n  +  2) 

It  will  be  noted  that  each  member  of  this  series  differs  from 
the  one  preceding  it  by  the  group  of  atoms  CH2,  and  that 
the  boiling  points  gradually  increase.  All  the  members  of 
this  series  are  known  up  to  the  one  having  the  formula 


204  FIRST  COURSE  IX  CHEMISTRY 

C28H5g.  The  lower  members  are  gases,  the  intermediate 
members  are  liquids,  and  the  higher  members  are  solids. 
They  are  all  combustible. 

Petroleum  and  products  derived  from  it.  This  liquid  is 
found  in  the  earth  in  certain  localities,  the  chief  oil-producing 
regions  in  the  United  States  being  Oklahoma,  California, 


FIG.  94.    Oil  wells  and  storage  tanks 

Illinois,  Texas,  and  Ohio.  By  means  of  compressed  air  or 
steam  the  petroleum  is  pumped  up  from  deep  wells  sunk 
into  the  ground  and  is  stored  in  large  tanks  (Fig.  94).  It 
is  composed  principally  of  liquid  hydrocarbons  in  which  are 
dissolved  both  gaseous  and  solid  hydrocarbons.  Crude  petro- 
leum is  not  only  used  as  fuel  in  this  country,  particularly 
on  locomotives  and  steamboats,  but  many  useful  products 
are  obtained  from  it  by  the  process  of  refining,  among  which 
are  gasoline,  kerosene,  lubricating  oils,  vaseline,  and  paraffin. 


CARBON  AND  ITS   SIMPLER  COMPOUNDS    205 

Refining  of  petroleum.  In  this  process  the  crude  oil  is  run  into 
large  iron  stills  (Fig.  95)  and  subjected  to  distillation.  The  dis- 
tillates which  pass  over  between  certain  limits  of  temperature 
are  kept  separate  and  serve  for  different  uses.  The  liquid 
passing  over  between  approximately  70°  and  150°  is  known  as 
naphtha,  while  that  passing  over  between  150°  and  300°  con- 
stitutes ordinary  kerosene.  A  number  of  different  naphthas  are 
recognized  commercially,  differing  in  boiling  points  and  density. 


FIG.  95.  Stills  for  refining  petroleum 

Those  of  low  boiling  point  constitute  ordinary  gasoline  and  are 
used  as  a  fuel  in  stoves  and  motors ;  those  of  higher  boiling 
points  are  used  in  making  paints.  Benzine  is  a  high-boiling 
naphtha,  and  being  a  good  solvent  for  such  organic  substances 
as  fats  and  oils,  is  used  in  cleaning  fabrics  (dry-cleaning). 

The  liquid  remaining  after  the  kerosene  and  higher-boiling 
oils  has  distilled  over  is  chilled,  whereupon  the  solid  constitu- 
ents dissolved  in  the  oil  separate.  These  are  filtered  off,  and 
constitute  ordinary  paraffin.  The  filtrate  is  then  distilled,  and 
from  it  various  lubricating  oils  are  obtained. 

Formerly  kerosene  was  the  most  important  of  the  products 
obtained  from  petroleum.  At  present,  however,  gasoline  is  the 


206 


FIKST  COURSE  IN  CHEMISTRY 


most  in  demand,  so  that  every  effort  is  made  to  increase  the 
yield  of  gasoline.  To  accomplish  this,  the  distillation  is  carried 
on  under  conditions  that  tend  to  decompose  the  heavier  mole- 
cules making  up  the  higher  boiling  liquids,  into  the  simpler 
molecules  which  constitute  liquids  of  lower  boiling  points.  The 

process  is  known  as  the  crack- 
ing of  oils. 

Methane  (CHJ.  This  hy- 
drocarbon, commonly  known 
as  marsh  gas,  is  formed  in 
marshes  and,  in  general, 
wherever  organic  matter  de- 
cays or  is  heated  in  the 
absence  of  air.  It  constitutes 
about  30  per  cent  of  coal  gas 
and  from  90  per  cent  to  95 
per  cent  of  natural  gas.  It 
often  collects  in  mines,  and 
when  mixed  with  air  is  called 
fire  damp.  Such  mixtures 


FIG.  96.    Sir  Humphry  Davy 
(1778-1829) 

A  distinguished  English  scientist,  who 


prepared    the   elements   sodium   and 
potassium 


invented   the   safety  lamp    and  first      are    very    explosive. 

Pure  methane  is  a  color- 
less, odorless  gas  about  one 
half  as  heavy  as  air.  It  is  but  slightly  soluble  in  water. 
When  ignited  it  burns  with  a  pale-blue  flame : 

CH4  +  2  O2  — >  CO2  +  2  H2O 

Safety  lamp.  Fortunately  the  ignition  point  of  fire  damp  (that 
is,  the  temperature  at  which  it  takes  fire)  is  high  and  its  flame 
may  be  extinguished  by  cooling.  In  1815  Sir  Humphry  Davy 
(Fig.  96)  invented  a  miner's  lamp  based  on  this  principle,  in 
which  the  usual  chimney  of  a  lantern  is  replaced  by  a  wire  gauze 
(Fig.  97).  An  explosion  flame  starting  at  the  wick  is  so  cooled 
by  the  metal  wire  that  ignition  ceases  and  the  explosion  is 


CARBON  AND  ITS   SIMPLER  COMPOUNDS    207 


FIG.  97.  Miner's  safety 
lamp 


confined  to  the  interior  of  the  lamp.  The  principle  may  be 
demonstrated  by  holding  a  wire  gauze  a  few  inches  above  a 
Bunsen  flame  and  parallel  with  the  table 
(Fig.  98).  When  the  gas  is  turned  on  and 
a  light  applied  above  the  gauze,  the  re- 
sulting flame  rests  upon  the  gauze,  but 
does  not  pass  through  it  to  the  burner. 

Halogen  derivatives  of  methane.  As  a  rule, 
the  hydrogen  present  in  a  hydrocarbon  may 
be  displaced  by  a  halogen  element,  atom 
for  atom.  In  this  way  there  are  formed 
from  methane  a  number  of  derivatives,  the 
most  important  of  which  are  the  following : 
Chloroform  (CHC13),  a  heavy,  colorless 
liquid  boiling  at  61°,  is  the  well-known 
anaesthetic  used  in  surgery.  Carbon  tetra- 
c/ilni-lde  (CC1.)  resembles  chloroform  in 

\  4/ 

appearance.    It  is  a  good  solvent,  espe- 
cially for  fatty  substances.    It  is  often 

used  to  remove  grease  spots  from  fabrics,  and  is  sold  for  this 
purpose  under  the  name  of  carbona.  It  possesses  the  advantage 
over  benzine  of  being  noninflammable,  but  is  more  expensive. 
lodoform  (CHIg)  is  a  yellow,  crystal- 
line solid  and  is  largely  used  as  an 
antiseptic  in  the  treatment  of  wounds. 

Acetylene  (C2H2).  This  is  a  color- 
less gas  having,  when  pure,  a  faint, 
pleasant  odor.  It  is  easily  obtained 
by  the  action  of  water  on  calcium 
carbide  (CaC2)  (p.  305) : 

CaC2  +  2  H20 >•  C2H2  +  Ca(OH)2 

In  this  way  the  gas  is  prepared  in  large 
quantities  for  use  as  an  illuminant  and  as  a  source  of  intense 
heat.  When  heated  it  decomposes  with  evolution  of  much 
heat :  C  H  ^  2  C  +  H  +  49,300  cal. 


FIG.  98.    An  experiment 

illustrating  the  principle 

of  the  safety  lamp 


208  FIKST  COURSE  IN  CHEMISTRY 

When  compressed  in  cylinders,  acetylene  is  very  explo- 
sive, since  the  heat  liberated  in  compressing  the  gas  is 
sufficient  to  start  decomposition.  With  the  proper  admix- 
ture of  air  it  burns  with  a  brilliant  white  light.  The  flame 
is  very  hot  because  to  the  heat  of  combustion  of  the  hy- 
drogen and  carbon  present  there  is  added  the  heat  of 
decomposition  of  the  acetylene  undergoing  combustion: 
2  C  H  +  5  O0 >-  4  CO,  +  2  H  O  +  603,260  cal. 

m     ,  m  m  i  2 


FIG.  99.    Cutting  an  iron  plate  by  means  of  the  oxyacetylene  blowpipe 

Uses  of  acetylene.  As  an  illuminant,  acetylene  is  chiefly 
used  when  electric  lights  are  not  available.  It  may  be 
safely  stored  in  metal  cylinders  by  filling  the  cylinder  with 
some  porous  material  (such  as  asbestos  and  cotton),  par- 
tially saturating  this  with  a  liquid  compound  called  acetone, 
and  then  forcing  in  the  gas  at  low  temperatures.  Under 
pressure,  the  acetone  dissolves  a  large  volume  of  the  gas. 
In  this  form  it  is  now  a  common  article  of  commerce. 


CARBON  AND  ITS  SIMPLER  COMPOUNDS    209 

The  intense  heat  generated  by.  the  combustion  of  acety- 
lene makes  it  useful  in  certain  processes  requiring  high 
temperatures,  such  as  the  welding  and  cutting  of  metals. 
For  this  purpose  the  acetylene  is  burned  in  an  apparatus 
known  as  the  oxyacetylene  blowpipe,  which  is  very  much 
like  the  oxyhydrogen  blowpipe.  A  temperature  of  about 
2700°  may  be  obtained  in  this  way.  This  blowpipe  has 
been-  found  especially  useful  in  cutting  iron  structures 
(Fig.  99),  since  the  tip  of  the  flame,  when  drawn  slowly 
over  the  metal,  melts  and  burns  it  at  the  point  of  contact, 
and  thus  makes  it  possible  to  cut  the  metal  into  pieces. 

EXERCISES 

1.  Suggest  a  method  for  proving  that  all  the  various  forms  of 
carbon  described  are  really  carbon. 

2.  Suggest  a  method  for  determining  the  percentage  of  carbon 
in  a  sample  of  coal. 

3.  How  could  you  distinguish  between  oxygen,  hydrogen,  nitro- 
gen, nitrous  oxide,  and  carbon  dioxide  ? 

4.  Why  cannot  coal  oil  be  substituted  for  gasoline  in  a  gaso- 
line engine  ? 

5.  Why  not  burn  gasoline  in  lamps? 

6.  Calculate  the  volumes  of  oxygen  necessary  to  burn  100  liters 
each  of  methane,  acetylene,  and  carbon  monoxide.  Ans.  200 ;  250 ;  50. 

7.  Suppose  that  gasoline  is  pure  heptane,  what  weight  of  oxy- 
gen would  be  necessary  to  burn  1  kg.  of  the  liquid?    Ans.  3515.5  g. 

TOPICS  FOR  THEMES 

(1)  Diamond  mines.  (2)  The  famous  diamonds.  (3)  Synthetic 
diamonds.  (For  information  on  diamonds,  refer  to  McPherson  and 
Henderson,  A  Course  in  General  Chemistry ;  Bird,  Modern  Science 
Reader ;  Duncan,  Chemistry  of  Commerce ;  and  encyclopedia.) 

Coal  mines  and  mining  (see  encyclopedia). 

The  production  and  refining  of  petroleum  (Rogers  and  Aubert, 
Industrial  Chemistry). 


CHAPTER  XXII 
FUELS  ;  FLAMES ;  ELECTRIC  FURNACES 

Fuels.  A  variety  of  substances  are  used  as  sources  of 
heat,  the  most  important  of  them  being  the  various  fuel 
gases,  together  with  coal,  wood,  and  petroleum.  The  com- 
position of  a  number  of  these  fuel  gases  is  given  in  the 
table  on  page  213.  Most  of  them  serve  as  illuminants  as 
well  as  fuels. 

Coal  gas.  It  has  been  known  for  several  centuries  that 
when  soft,  or  bituminous,  coal  is  heated  out  of  contact  with 
air,  combustible  gases  are  formed ;  indeed,  gas  obtained  in 
this  way  was  used  for  street  lighting  in  London  and  Paris 
a  hundred  years  ago. 

The  manufacture  of  coal  gas.  The  manufacture  of  coal  gas  is 
represented  in  a  diagrammatic  way  in  Fig.  100.  The  coal  is 
introduced  into  a  closed  retort  J,  and  heated  by  the  fire  below. 
A  number  of  these  retorts  are  placed  in  parallel  rows,  each 
being  furnished  with  a  delivery  pipe,  from  which  the  gas 
bubbles  into  the  tarry  liquids  which  collect  in  the  hydraulic 
main  B,  running  above  the  retorts.  In  this  large  pipe  are  de- 
posited most  of  the  solid  and  liquid  products  formed  in  distil- 
lation, constituting  the  sticky  mass  known  as  coal  tar.  On  the 
top  of  the  tar  there  collects  a  liquid  containing  ammonia  and 
known  as  the  ammoniacal  li'/nor.  The  partially  purified  gas 
then  passes  into  a  series  of  pipes  C,  in  which  it  is  cooled  and 
further  separated  from  tar.  In  the  scrubber  D  it  passes 
through  a  column  of  loose  coke,  over  which  water  is  sprayed, 
where  it  is  still  further  cooled  and  to  some  extent  purified 

210 


FUELS;    FLAMES;    ELECTRIC  FURNACES     211 

from  soluble  gases,  such  as  hydrogen  sulfide  and  ammonia.  In 
the  purifier  E  it  passes  over  a  bed  of  lime  or  of  iron  oxide, 
which  removes  the  remainder  of  the  sulfur  compounds,  and 
from  this  it  enters  the  large  gas  holder  Ft  from  which  it  is 
distributed  to  consumers. 

The  great  bulk  of  the  carbon  remains  in  the  retort  as  coke 
and  as  retort  carbon.  The  yield  of  gas,  tar,  and  soluble  materials 
depends  upon  many  factors,  such  as  the  composition  of  the 
coal,  the  temperature  employed,  and  the  rate  of  heating.  One 


FIG.  100.    Plant  for  the  manufacture  of  coal  gas  and  accompanying 
products 

ton  of  good  gas  coal  yields  approximately  10,000  cu.  ft.  of  gas, 
1400  Ib.  of  coke,  120  Ib.  of  tar,  and  20  gal.  of  ammoniacal  liquor. 
Not  only  is  the  ammonia  obtained  in  the  manufacture  of  the 
gas  of  great  importance,  but  the  coal  tar  is  the  source  of  many 
very  useful  substances,  as  will  be  explained  later. 

Water  gas.  Water  gas  is  essentially  a  mixture  of  carbon 
monoxide  and  hydrogen.  It  is  made  by  passing  steam  over 
very  hot  anthracite  coal  or  coke,  when  the  reaction  shown 
in  the  following  equation  takes  place : 

C  +  H20  — >-  CO  +  H2 

The  process  is  intermittent.    The  fuel  is  burned  in  a  draft 
of  air  until  it  is  very  hot.    The  air  is  then  shut  off  and 


212  FIRST  COURSE  IN  CHEMISTRY 

steam  turned  on.    The   temperature   gradually   falls,   and 
when  it  reaches  about  1000°,  the  process  is  again  reversed. 

Water  gas  is  very  effective  as  a  fuel,  since  both  carbon 
monoxide  and  hydrogen  burn  with  very  hot  flames.  It  has 
little  odor  and  is  very  poisonous.  Its  use  is  therefore  at- 
tended with  some  risk,  since  leaks  in  pipes  are  very  likely 
to  escape  notice. 

Enriched  water  gas.  When  required  merely  for  the  produc- 
tion of  heat,  the  gas  as  prepared  above  is  at  once  ready  for  use. 
When  made  for  illuminating  purposes,  it  must  be  enriched ; 
that  is,  illuminants  must  be  added,  since  both  carbon  monoxide 
and  hydrogen  burn  with  a  nonluminous  flame.  This  is  accom- 
plished by  passing  it  into  heaters  containing  highly  heated 
petroleum  oils.  The  gas  takes  up  the  gaseous  hydrocarbons 
formed  in  the  decomposition  of  the  petroleum  oils,  and  these 
hydrocarbons  make  it  burn  with  a  luminous  flame. 

Producer  gas.  This  is  made  by  burning  coal  in  a  limited 
supply  of  air  so  that  the  product  of  combustion  is  largely 
carbon  monoxide.  It  is  used  as  a  fuel  in  gas  engines  and 
in  many  industrial  operations. 

Natural  gas.  In  many  regions  of  the  United  States,  as 
well  as  in  other  countries,  natural  gas  is  obtained  from 
wells  drilled  into  a  stratum  holding  the  gas.  While  it  is 
variable  in  composition,  it  consists  largely  of  methane,  many 
samples  containing  as  much  as  95  per  cent  of  this  com- 
pound. It  burns  with  a  rather  smoky  flame  of  moderate 
luminosity,  but  works  well  with  a  gas  mantle.  It  has  a  high 
heat  of  combustion,  as  shown  in  the  following  equation : 

CH4  +  2  02 >  C02  +  2  H20  +  213,500  cal. 

It  is  an  ideal  fuel  and  is  often  conducted  through  pipes 
for  hundreds  of  miles  from  the  gas  fields  to  cities. 


FUELS;   FLAMES;   ELECTRIC  FUKNACES     213 


Comparative  composition  of  fuel  gases.  The  following  figures 
are  the  results  of  analyses  of  average  samples,  but  since  each 
kind  of  fuel  gas  varies  considerably  in  composition,  the  values 
are  to  be  taken  as  approximate  only.  The  nitrogen  and  traces 
of  oxygen  are  derived  from  the  air. 


COMPOSITION  OF  GASES 


CONSTITUENT 

OHIO 
NATURAL 
GAS 

COAL, 
GAS 

WATER 

GAS 

ENRICHED 
WATER 
GAS 

PRODUCER 

GAS 

II 

09 

41  3 

5288 

3796 

10  90 

C  H6       

89.5 
9.3 

43.6^ 

2.16 

7.09 
2.01 

C,II2  and  C2H4     .     . 
CO     
CO 

0.3 
0.4 
0.3 

3.9 
6.4 
2.0 

36.80 
3.47 

9.40 
32.25 
4.73 

0.60 
20.10 

8.50 

0.2 

0.0 

1.2 
0.3 

4.69 

3.96 
0.60 

59.90 

Other  hydrocarbons  . 

0.0 

1.5 

1.80 

Gas  mantles.  In  using  the  fuel  gases  as  illuminants,  the 
gas  is  generally  mixed  with  air  before  burning.  In  this  way 
the  gas  burns  with  a  hot  but  nearly  nonluminous  flame. 
The  light  is  obtained  by  suspending  about  this  flame  a 
gauze  mantle  of  suitable  material.  The  best  mantles  are 
composed  of  a  mixture  of  99  per  cent  of  thorium  oxide 
with  1  per  cent  of  cerium  oxide. 

The  thorium  and  cerium  compounds  used  in  gas  mantles 
are  obtained  from  monazite  sand  (Fig.  101)  found  prin- 
cipally in  North  Carolina  and  Brazil.  The  process  of  mak- 
ing a  gas  mantle  consists  in  knitting  a  cylindrical  cotton 
fabric,  which  is  then  dipped  into  a  solution  of  the  nitrates 
of  thorium  and  cerium.  After  drying,  the  fabric  is  heated, 
in  which  process  the  yarn  is  burned,  while  the  nitrates 
of  thorium  and  cerium  are  converted  into  oxides  (p.  131) 


214 


FIRST  COURSE  IN  CHEMISTRY 


which  are  left  in  the  form  of  the  original  fabric.  The  result- 
ing mantle  is  very  delicate,  and  is  strengthened  for  shipping 
by  dipping  it  into  a  solution  of  an  appropriate  substance 
and  drying. 

Products  of  the  combustion  of  ordinary  fuels.  Ordinary 
fuels,  such  as  oil,  wood,  coal,  and  fuel  gases,  are  largely 
made  up  of  carbon  and  hydrogen  or  their  compounds.  The 
chief  products  of  the  combustion  of  such  fuels  are  carbon 


FIG.  101.    Materials  used  in  making  gas  mantles ;  also  different 
stages  in  the  process 

dioxide  and  water.  Associated  with  these  are  small  amounts 
of  other  products,  such  as  carbon  monoxide  and  sulfur 
dioxide,  the  latter  being  formed  from  traces  of  sulfur 
compounds  in  the  fuels. 

Rooms  are  not  infrequently  heated  by  gas  or  oil  stoves, 
with  no  provisions  for  removing  the  products  of  combustion. 
Likewise,  natural  gas  is  often  burned  in  stoves  or  grates 
with  the  damper  closed  so  as  to  leave  no  opening  into  the 
chimney.  Such  practices  are  greatly  to  be  condemned,  since 


FUELS;   FLAMES;   ELECTRIC  FURNACES     215 


the  air  in  the  rooms  heated  in  this  way  soon  becomes  so 
contaminated  with  the  various  products  of  combustion  as 
to  render  it  unfit  for  respiration.  The  large  amount  of 
water  vapor  formed  in  rooms  so  heated  condenses  on  the 
windows  in  cold  weather,  causing  the  glass  to  siveat. 

The  electric  furnace.  In  recent  years  electric  furnaces  have 
come  into  wide  use  in  operations  requiring  a  very  high  tem- 
perature. Temperatures  as  high  as  3500°  can  easily  be  reached. 
These  furnaces  are  constructed  on 
one  of  two  general  principles. 

1.  Arc  furnaces.   In  the  one  type 
the  source  of  heat  is  an  electric  arc 
formed  between  carbon  electrodes 
separated  a  little  from  each  other, 
as   shown   in  Fig.  102.     The   sub- 
stance to  be  heated  is  placed  in  a 
vessel,  usually  a  graphite  crucible, 
just  below  the  arc.    The  electrodes 

and  crucible  are  surrounded  by  materials  which  fuse  with 
great  difficulty,  such  as  magnesium  oxide,  the  walls  of  the 
furnace  being  so  shaped  as  to  reflect  the  heat  downwards  upon 
the  contents  of  the  crucible. 

2.  Resistance  furnaces.   In  the  other  type  of  furnace  the  heat  is 
generated  by  the  resistance  offered  to  the  current  in  its  passage 
through  the  furnace.   A  typical  form  of  such  a  furnace  is  illus- 
trated in  Fig.  90,  which  is  used  in  the  manufacture  of  graphite. 

Conditions  necessary  for  flames.  When  one  of  the  sub- 
stances undergoing  combustion  remains  solid  at  the  tem- 
perature occasioned  by  the  combustion,  light  may  be  given 
off,  but  there  is  no  flame.  Thus,  iron  wire  burning  in 
oxygen  throws  off  a  shower  of  sparks,  but  no  flame  is  seen. 
When,  however,  both  of  the  substances  involved  are  gases 
or  vapors  at  the  temperature  reached  in  the  combustion, 
the  act  of  union  is  accompanied  by  a  flame. 


FIG.  102.    Electric  furnace  of 
the  arc  type 


216 


FIKST  COURSE 


CHEMISTEY 


Flames  from  burning  liquids  or  solids.  Many  substances 
which  are  liquids  or  solids  at  ordinary  temperatures  burn 
with  a  flame  because  the  heat  of  combustion  slowly  vaporizes 
them,  and  the  flame  is  due  to  the  union  of  this  vapor  with  the 
oxygen  of  the  air.  This  may  be  shown  in  the  case  of  a  candle 

flame  by  holding  one  end  of  a 
slender  glass  tube  in  the  base 
of  the  flame  (Fig.  103).  The  un- 
burned  vapor  in  the  inner  part 
of  the  flame  is  thus  conducted 
away,  and  may  be  ignited  at  the 
upper  end  of  the  tube. 

Structure  of  a  flame.  When 
hydrogen  or  carbon  monoxide 
burns  in  oxygen, 
but  one  reaction 
takes  place,  and 
as  a  result 
the  flame  is 

very  simple  in  structure.    It  consists  of  a 

colorless  inner  cone  of  unburned  gas  and 

an  outer  cone  in  which  the  union  between 

the  hydrogen  and  oxygen  is  taking  place. 

It  follows  that  the  outer  cone  is  the  hot 

part  of  the  flame.    That  the  inner  cone  is 

cool  is  shown  by  the  fact  that  a  match 

head  suspended  in  this  region  (Fig.  104) 

before   lighting   the   gas,  and   left  there 

while  the  gas  burns,  is  not  ignited. 
The  flames  produced  by  the  combustion 

of  hydrocarbons  such  as  are  present  in 

coal  gas  and  natural  gas  or  of  mixtures  of 

hydrocarbons  with  stearic  acid,  as  in  candles,  is  much  more 

complex  because  several  consecutive  reactions  take  place. 


FIG.  103.    Method  of  proving  that 
the  interior  of  a  candle  flame  con- 
tains combustible  vapors 


FIG.  104.  A  match 
head  suspended  in 
lower  part  of  gas 
flame  is  not  ignited 


FUELS;    FLAMES;    ELECTEIC  FUENACES     217 


FIG.  105.   The  cones  of 
a  candle  flame 


For  example,  in  the  candle  flame  (Fig.  105)  there  are, 
broadly  speaking,  three  cones:  (1)  the  inner  cone  A,  com- 
posed of  combustible  vapors ;  (2)  an 
intermediate  cone  B,  in  which  these 
vapors  are  decomposed  by  the  heat 
and  a  small  quantity  of  carbon  is  set 
free  which  renders  the  flame  luminous ; 
and  (3)  an  almost  invisible,  narrow 
outer  cone,  or  film,  (7,  in  which  the 
carbon  and  hydrogen  are  burned  to 
water  and  carbon  dioxide. 

Bunsen  burners.  In  the  ordinary 
Bunsen  burner,  and  in  similar  burners 
used  in  gas  ranges  (Fig.  106)  and  for 
illumination  with  the  aid  of  mantles, 
the  gas  is  mixed  with  a  certain  per- 
centage of  air  before  it  is  burned. 

This  is  accomplished  by  having  an  opening  (mixer)  in  the 
base  of  the  burner  (Fig.  106,  A)  into  which  the  air  is  drawn 
by  the  flow  of  the  gas.  If  the 
mixer  is  adjusted  so  that  the 
proper  amount  of  air  is  ad- 
mitted, the  flame  is  color- 
less. Such  a  flame  possesses 
an  advantage  in  that  it  is 
very  hot  and  no  carbon  is 
deposited  from  it. 

Smoke   prevention.    Since 
the  products  of  combustion 
of  fuels  are  carbon  dioxide       FlG<  106>  A  typical  gas  burner 
and  water  vapor,  and  these 

are  invisible  compounds,  it  is  evident  that  if  the  combus- 
tion is  complete,  no  smoke  will  be  formed.   As  a  rule  the 


218 


FIRST  COURSE  IN  CHEMISTRY 


combustion  is  imperfect ;  gaseous  compounds  containing 
carbon  are  first  formed,  and  when  these  are  imperfectly 
burned,  a  part  of  their  carbon  is  set  free  in  a  finely  divided 
state  constituting  smoke.  Smoke  may  therefore  be  pre- 
vented by  securing  the  complete  combustion  of  the  fuel, 
the  necessary  conditions  being  as  follows:  (1)  a  sufficient 


|f: 
FIG.  107.  Diagram  of  a  smoke  consumer 

supply  of  air;  (2)  thorough  mixing  of  the  air  with  the 
combustible  gases  produced  from  the  fuel ;  and  (3)  a  tem- 
perature high  enough  to  maintain  active  combustion. 

Smoke  prevention  is  a  problem  of  great  economic  importance, 
especially  in  the  large  cities.  Thus,  for  example,  it  is  estimated 
that  the  smoke  in  the  city  of  Pittsburgh  costs  the  people  of 
the  city  $10,000,000  yearly,  or  about  $20  for  each  inhabitant ; 
and  this  does  not  take  into  account  the  serious  effect  of  smoke 
upon  health.  Because  of  these  facts  many  cities  are  now 
taking  steps  to  abate  the  smoke  nuisance.  That  the  conditions 


FUELS;   FLAMES;   ELECTRIC  FURNACES     219 

necessary  for  preventing  smoke  may  be  met,  it  is  essential  that 
the  coal  be  introduced  into  the  furnace  uniformly,  so  that  the 
volatile  matter  expelled  upon  heating  may  be  more  readily 
mixed  with  air.  This  is  done  efficiently  by  having  a  chain 
grate,  as  is  shown  in  A  (Fig.  107).  The  coal  is  fed  on  this  at 
B,  and  as  the  chain  slowly  moves  forward,  the  coal  gradually 
enters  the  furnace,  and  by  the  time  it  reaches  the  back  part  of 
the  furnace,  C,  it  is  completely  burned,  the  ashes  falling  out 
at  D.  The  volatile  matter  expelled  is  thoroughly  mixed  with 
hot  air  led  in  through  the  back  of  the  grate  E,  E.  The  large 
space  under  the  boiler  drum  gives  opportunity  for  complete 
combustion  of  the  products  under  the  chimney.  The  water  in 
the  drum  circulates  through  the  tubes,  as  shown  by  the  arrows, 
and  thus  is  heated  to  a  high  temperature. 

EXERCISES 

1.  Why  does  charcoal  usually  burn  with  no  flame?    How  do  you 
account  for  the  flame  sometimes  observed  when  it  burns  ? 

2.  Would  anthracite   coal  be   suitable   for  the   manufacture  of 
coal  gas  ? 

3.  Suggest  a  way  in  which  natural  gas  may  have  been  formed. 

4.  Why  does  the  use  of  the  bellows  on  the  blacksmith's  forge 
cause  a  more  intense  heat  ? 

5.  Assuming  that  natural  gas  is  composed  wholly  of  marsh  gas, 
what  will  be  the  weight  of  1  cu.  m.  of  it  (standard  conditions)? 
Ans.    716.8  g.    What  will  be   the  weight  of  water   and  of  carbon 
dioxide  formed  in  its  combustion?    Ans.  1611  g. ;  1967  g. 

6.  A  portable  gas  stove  used  in  heating  a  room  burns  10  cu.  ft. 
of  gas  each  hour.    Supposing  that  the  gas  is  pure  methane,  what 
volume   of   oxygen  is  withdrawn  each   hour  from  the   air  in  the 
room  ?    A  ns.  20  cu.  ft.  What  volume  of  carbon  dioxide  is  given  off  ? 
Ans.    10  cu.  ft. 

TOPICS  FOR  THEMES 

The  production  and  uses  of  natural  gas.    (Write  to  your  state 
geologist  for  references.) 

The  use  of  high  temperatures  (Duncan,  Chemistry  of  Commerce). 
Methods  of  heating.    (Consult  local  dealers  in  stoves  and  furnaces.) 


CHAPTER  XXIII 
CARBOHYDRATES;    ALCOHOLS;    COAL-TAR  COMPOUNDS 

Carbohydrates.  The  term  carbohydrate  is  applied  to  a 
class  of  compounds  which  includes  the  sugars,  starch,  and 
allied  substances.  These  compounds  contain  carbon,  hydro- 
gen, and  oxygen,  the  last  two  elements  usually  being  pres- 
ent in  the  proportion  in  which  they  combine  to  form  water. 
The  most  important  carbohydrates  are  the  following : 

TABLE  OF  CARBOHYDRATES 

Sucrose  (ordinary  sugar) ^12^22^11 

Lactose  (milk  sugar) C12H(>2H11 

Maltose C12H22On 

Dextrose  (grape  sugar) C6H12O6 

Levulose C6H12O6 

Cellulose (C6H10O5)X 

Starch (C6H10O5)X 

The  molecular  formulas  of  cellulose  and  starch  are  unknown 
but  are  multiples  of  the  simple  formula  C6H1QO5;  hence 
they  are  often  written  (C6H10O5)X.  In  the  discussion  of  the 
compounds  they  will  be  represented  by  the  simple  formula 

C6H1005- 

It  will  be  noted  that  some  of  the  compounds  named  in 

the  above  table  have  the  same  formula.  Compounds  having 
the  same  formula  are  said  to  be  isomeric.  The  difference 
in  the  properties  of  such  compounds  is  due  to  the  fact  that 
the  atoms  are  arranged  differently  in  the  molecule. 

220 


CARBOHYDRATES 


221 


Sucrose  (sugar)  (C12H22On).  This  substance,  commonly 
called  sugar,  occurs  in  many  plants,  especially  in  the  sugar 
cane  and  sugar  beet,  each  of  which  at  present  furnishes 
about  50  per  cent  of  the  total  production.  The  sugar  cane 
grows  only  in  warm  climates  (Cuba  and  the  Hawaiian 
Islands  are  the  greatest  producers),  while  the  sugar  beet 
thrives  in  cooler  climates,  such 
as  prevail  in  Ohio  and  Michi- 
gan in  the  United  States,  and  | 
in  Germany.  The  beets  con- 
tain as  high  as  16  per  cent  of 
sucrose. 

The  manufacture  of  sugar.  The 
juice  from  the  cane  or  beet  con- 
tains the  sugar  in  solution  along 
with  many  impurities.  These  im- 
purities are  removed,  partly  by 
precipitation  and  partly  by  filter- 
ing through  bone  black,  and  the 
resulting  solution  is  then  evapo- 
rated until  the  sugar  crystallizes. 
The  evaporation  is  conducted  in 
closed  vessels  from  which  the  air 
is  partially  exhausted  (vacuum 

pans  (Fig.  108)).  In  this  way  the  boiling  point  of  the  solution 
is  lowered  and  the  charring  of  the  sugar  is  prevented.  It  is 
not  practical  to  remove  all  the  sugar  from  the  solution.  Ordi- 
nary molasses  is  the  solution  which  remains  after  the  sugar 
has  been  crystallized  out  from  the  purified  juice  of  the  sugar 
cane.  The  sweetness  of  maple  sugar  is  due  to  sucrose,  other 
products  present  in  the  maple  sap  imparting  the  distinctive 
flavor.  About  40,000,000,000  Ib.  of  sugar  is  produced  annu- 
ally. The  annual  consumption  of  sugar  in  the  United  States 
amounts  to  about  8,250,000,000  Ib.  or  approximately  85  Ib.  for 
each  person. 


FIG.  108.    Vacuum  pans  in  a 
sugar  factory 


222  FIRST  COURSE  IN  CHEMISTEY 

Chemical  conduct  of  sugar.  When  a  solution  of  cane 
sugar  is  heated  to  about  70°  with  hydrochloric  acid,  two 
isomeric  sugars,  dextrose  and  levulose,  are  formed  in  accord- 
ance with  the  following  equation : 

C]2H22On  +  H20  -*•  CCHJ30S  +  C6H1206 

In  this  process  the  sugar  is  said  to  be  inverted,  and  the 
mixture  of  dextrose  and  levulose  is  termed  invert  sugar. 

When  heated  to  160°,  sucrose  melts ;  if  the  temperature  is 
increased  to  about  215°,  a  partial  decomposition  takes  place, 
and  a  brown  substance  known  as  caramel  is  formed.  This 
is  used  extensively  as  a  coloring  matter  and  in  making 
confectionery. 

Lactose  (milk  sugar)  (C12H22011).  This  compound  is  present 
in  the  milk  of  all  mammals.  The  average  composition  of 
cow's  milk  is  as  follows: 

Water 87.17% 

Casein  (nitrogenous  matter) 3.56% 

Butter  fat 3.64% 

Lactose 4.88% 

Mineral  matter 0.75% 

When  rennin  (a  substance  obtained  from  the  stomach  of 
calves)  is  added  to  milk,  the  casein  separates  and  is  used  in 
the  manufacture  of  cheese.  The  remaining  liquid,  known 
as  whey,  contains  the  milk  sugar,  which  crystallizes  on  evap- 
oration ;  it  resembles  sucrose  in  appearance,  but  is  not  so 
sweet  or  so  soluble.  The  souring  of  milk  is  due  to  the  fact 
that  the  milk  sugar  contained  in  it  changes  into  lactic  acid, 
a  liquid  having  the  formula  C3H6O3 : 

CBHBOn  +  HfO— H4C.H.O, 

This  change  is  brought  about  through  the  agency  of  a 
certain  microorganism  which  enters  from  the  air,  and  the 
process  is  known  as  lactic  fermentation.  The  body  of  the 


CARBOHYDRATES  223 

ordinary  medicine  tablet  consists  of  lactose,  because  this 
substance  readily  absorbs  medicinal  solutions. 

Dextrose  (grape  sugar,  glucose)  (C6H1206).  This  sugar 
is  present  in  honey  and  in  many  fruits,  usually  associated 
with  levulose,  and  is  often  called  grape  sugar  because  of  its 
presence  in  grape  juice.  It  can  be  obtained  along  with 
levulose  by  heating  sucrose  with  hydrochloric  acid,  as  ex- 
plained above.  Commercially  it  is  prepared  in  enormous 
quantities  by  heating  starch  with  hydrochloric  acid.  The 
starch  is  first  changed  into  a  sweet-tasting  solid  known  as 
dextrin,  and  this,  on  further  action,  is  converted  into  dextrose : 

C.HWO.+H,O— »-C.HMO, 

When  the  change  is  complete,  the  hydrochloric  acid  is 
neutralized  by  sodium  carbonate.  Over  50,000,000  bushels 
of  corn  are  used  each  year  in  the  United  States  in  the 
production  of  dextrose  and  allied  products. 

Pure  dextrose  is  a  white,  crystalline  solid  resembling 
sucrose  in  its  properties,  but  is  not  so  sweet.  Most  of 
the  dextrose  used  is  in  the  form  known  commercially  as 
glucose,  or  corn  sirup.  This  is  a  thick,  sirupy  liquid  and 
contains,  in  addition  to  water,  from  30  per  cent  to  40  per 
cent  dextrose,  and  from  40  per  cent  to  50  per  cent  dextrin. 
Large  quantities  of  glucose  are  used  in  the  preparation  of 
jellies,  jams,  sirups,  candy,  and  other  sweets.  The  federal 
law  requires  that  when  glucose  is  present  in  such  foods 
as  jellies  and  jams,  the  label  on  the  container  must  state 
the  percentage  of  glucose  present. 

Starch  (C6H1005).  This  substance  is  always  present  in 
seeds  and  tubers  and  is  by  far  the  most  abundant  carbo- 
hydrate found  in  nature.  In  the  United  States  it  is  obtained 
chiefly  from  corn,  about  60  per  cent  of  which  is  starch.  In 
Europe  the  potato  serves  as  the  principal  source. 


224 


FIRST  COURSE  IN  CHEMISTRY 


The  manufacture  of  starch.  In  manufacturing  starch  from  corn, 
the  corn  is  first  soaked  in  water  containing  a  little  sulfurous 
acid,  to  soften  the  grain.  It  is  then  ground  coarsely  so  as  not 
to  crush  the  germ.  When  the  resulting  mass  is  mixed  with 
water,  the  germ  floats,  being  very  light  because  of  the  oil  which 
it  contains.  In  this  way  the  germ  is  separated  from  the  rest 
of  the  seed,  and  from  it  corn  oil  is  prepared.  The  remaining 
material,  consisting  of  the  starch,  the  nitrogenous  constituent 
(gluten),  and  the  bran,  or  outside  coating  of  the  grain,  is  then 


FIG.  109.   The  interior  of  a  starch  factory,  showing  the  settling  troughs 

ground  finely,  mixed  with  water,  and  passed  through  cloth 
sieves,  which  remove  the  bran.  The  water  containing  the 
starch  and  gluten  in  suspension  is  then  allowed  to  run  slowly 
down  long,  shallow  troughs,  the  rate  of  flow  being  regulated  so 
that  the  heavier  starch  sinks  to  the  bottom  of  the  trough  while 
the  lighter  gluten  is  washed  away.  The  starch  is  then  removed 
from  the  troughs  (Fig.  109)  and  dried.  Large  quantities  of 
starch  are  used  in  making  glucose  and  other  foods,  for  finishing 
cloth,  and  for  laundry  purposes. 

Characteristics  of  starch.  Starch  consists  of  minute  gran- 
ules, which  differ  somewhat  in  appearance,  according  to 
the  source  of  the  starch,  so  that  it  is  often  possible  from  a 


CARBOHYDBATES 


225 


microscopic  examination  to  determine  from  what  plant  any 
given  sample  of  starch  was  obtained  (Figs.  110  and  111). 
When  heated  with  water,  the  granules  burst,  and  the 
starch  partially  dissolves.  This  is  the 
reason  why  starchy  foods  are  made 
more  digestible  by  cooking. 

Cellulose  (C6H1005).  This  forms  the 
basis  of  all  woody  fibers.  Cotton  and 
linen  are  nearly  pure  cellulose.  It  is 
insoluble  in  water,  alcohol,  and  dilute 
FIG.  110.  Cornstarch  acids,  but  Avill  dissolve  in  a  solution 
gTanules^magnified  200  preparecl  by  dissolving  copper  oxide  ill 
ammonium  hydroxide.  Sulfuric  acid 
slowly  changes  it  into  dextrose.  Concentrated  nitric  acid 
forms  a  mixture  of  compounds  commonly  known  as  nitro- 
cellulose or  guncotton.  These  are  very  inflammable  and  under 
certain  conditions  are  highly  explosive.  They  have  many 
commercial  uses.  Photographic  films  are  made  from  them, 
as  well  as  from  a  noninflammable  derivative  of  cellulose 
known  as  acetyl  cellulose.  Collo- 
dion is  a  solution  of  nitrocellu- 
lose in  a  mixture  of  alcohol  and 
ether.  Celluloid  is  a  mixture 
of  nitrocellulose  and  camphor. 
These  two,  when  mixed  together, 
form  a  plastic  mass  which  can  be 
molded  into  any  desired  shape 
and  which  is  used  for  making 
such  objects  as  combs  and  brush 
handles. 


FIG.  111.    Wheat-starch  gran- 
ules magnified  200  diameters 


Mercerized  cotton  and  artificial  silk.  When  cotton  cloth  is 
treated  with  a  concentrated  solution  of  sodium  hydroxide,  the 
cellulose  shrinks  and  becomes  tougher  in  character.  If  the  cloth 


226  FIRST  COURSE  IN  CHEMISTRY 

is  placed  in  stretchers  to  prevent  the  shrinkage,  it  assumes  an 
appearance  somewhat  resembling  silk,  and  is  known  as  mercer- 
ized cotton.  Another  fabric  prepared  in  large  quantities  from 
cellulose  resembles  silk  very  closely,  and  is  known  as  artificial 
silk.  The  fiber  of  this  fabric  is  prepared  by  forcing  concen- 
trated solutions  of  cellulose  or  its  derivatives  through  minute 
tubes  and  coagulating  the  cellulose  as  it  emerges  in  the  form 
of  fine  threads. 

Characteristics  of  various  textile  fibers.  Of  the  different 
fibers  used  in  making  the  yarns  from  which  the  common 
fabrics  are  prepared,  the  vegetable  fibers,  cotton  and  linen, 


Silk  liber  Cotton  fiber  Wool  fiber 

FIG.  112.    Textile  fibers 

are  essentially  cellulose,  while  the  animal  fibers,  wool  and 
silk,  are  composed  of  nitrogenous  substances.  Although 
these  fibers  resemble  each  other  when  viewed  with  the  naked 
eye,  their  appearance  varies  widely  when  examined  with 
the  microscope.  The  characteristic  appearance  of  these 
fibers  is  shown  in  Fig.  112.  It  is  also  possible  to  distin- 
guish between  the  fibers  by  the  action  of  chemical  reagents. 
For  example,  a  hot  solution  of  sodium  hydroxide  (5  per  cent 
to  10  per  cent)  has  but  little  action  upon  cotton,  while  it  will 
readily  dissolve  wool  and  slowly  dissolve  silk. 

Paper.  Paper  consists  mainly  of  cellulose,  the  finer  grades 
being  made  from  linen  and  cotton  rags  and  the  cheaper 
grades  from  wood. 


CABBOHYDBATES  227 

Manufacture  of  paper.  In  making  paper,  the  raw  material  is 
cut  into  pieces  and  treated  with  suitable  reagents  (calcium 
bisulfite  is  used  in  case  of  wood),  to  remove  all  objectionable 
matter,  leaving  the  cellulose,  which  is  then  bleached  with  chlo- 
rine. The  paper  pulp  so  obtained  is  suspended  in  water  and  run 
onto  wire  screens.  It  then  passes  between  large  iron  cylinders, 
some  of  which  are  heated  with  steam.  In  this  way  the  pulp  is 
pressed  and  dried  and  delivered  in  the  form  of  paper.  In  the 


FIG.  113.   The  interior  of  a  paper  mill 

process  different  materials  are  often  added  to  the  pulp.  These 
vary  with  the  nature  of  the  paper  desired ;  thus,  finely  ground 
clay  or  calcium  sulfate  is  added  to  give  body  to  the  paper.  In 
making  paper  intended  for  writing  or  printing,  a  compound  pre- 
pared by  heating  resin  and  sodium  hydroxide  is  added,  together 
with  aluminium  sulfate.  This  prevents  the  ink  from  spreading. 
Fig.  113  shows  the  interior  of  the  paper  mill.  The  pulp 
flows  from  the  container,  A,  onto  the  screens  beyond  and  then 
between  the  rollers  until  it  is  pressed  and  dried  and  so  converted 
into  the  finished  paper,  B. 


228  FIBST  COUBSE  IN  CHEMISTBY 

The  alcohols.  A  great  many  alcohols  are  known.  The 
two  most  important  ones  are  methyl  alcohol  (CH3OH)  and 
ethyl  alcohol  (C2H5OH). 

Methyl  alcohol  (wood  alcohol)  (CH3OH).  This  compound 
is  formed  when  wood  is  heated  in  the  absence  of  air 
(p.  196),  and  on  this  account  it  is  called  wood  alcohol  It 
is  a  colorless  liquid  which  boils  at  64.7°  and  burns  with  an 
almost  colorless  flame.  It  is  a  good  solvent  for  organic  sub- 
stances and  is  used  extensively  in  the  manufacture  of  var- 
nishes. It  is  quite  poisonous.  It  acts  upon  the  optic  nerve, 
and  many  persons  have  become  blind  from  drinking  the 
liquid  or  from  repeatedly  inhaling  its  vapor. 

When  a  mixture  of  the  vapor  of  methyl  alcohol  and  air  is 
passed  over  hot  copper,  the  alcohol  is  oxidized,  forming  a 
gaseous  compound  known  as  formaldehyde : 

2  CH3OH  +  02 >•  2  CH20  +  2  H20 

This  gas  is  now  prepared  in  large  quantities  and  used  as  a 
disinfectant.  A  40  per  cent  aqueous  solution  of  it  is  sold 
under  the  name  of  formalin. 

Ethyl  alcohol  (grain  alcohol,  alcohol)  (C2H5OH).  This 
compound  is  the  one  commonly  known  as  alcohol.  It 
resembles  methyl  alcohol  in  its  general  properties. 

1.  Preparation.  It  is  prepared  by  the  action  of  ordinary 
baker's  yeast  upon  different  sugars  such  as  dextrose : 

C6H1206  — >•  2  C2H6OH  +  2  CO2 

This  process  in  which  a  sugar  is  changed  into  alcohol  and 
carbon  dioxide  by  the  action  of  yeast  is  known  as  alcoholic 
fermentation.  The  yeast  is  a  low  form  of  plant  life  (Fig.  114), 
and  thrives  in  appropriate  sugar  solutions.  During  its  growth 
a  number  of  changes  take  place  which  result  in  converting 
the  sugar  into  alcohol. 


ALCOHOLS 


229 


FIG.  114.    Some  cells  of  the 
yeast  plant 


Experimental  preparation  of  alcohol.  The  formation  of  alcohol 
and  carbon  dioxide  from  dextrose  may  be  shown  as  follows : 
A  10  per  cent  solution  of  the  sugar  in  water  is  poured  into 

flask  A  (Fig.  115)  and  a  little  baker's 
yeast  is  added.  The  bottle  B  con- 
taining limewater  is  connected  as 
shown  in  the  figure.  The  tube  C  is 
filled  with  pieces  of  sodium  hy- 
droxide. The  temperature  is  main- 
tained at  about  30°.  Action  soon 
begins,  as  is  indicated  by  the  rising 
bubbles  of  carbon  dioxide,  and  con- 
tinues for  some  hours  until  the  sugar 
is  all  fermented.  That  the  escaping 
gas  is  carbon  dioxide  is  shown  by 
the  precipitate  formed  in  B.  The 
sodium  hydroxide  in  C  prevents  carbon  dioxide  from  entering 
from  the  air.  The  alcohol  formed  is  separated  by  distillation. 
Commercial  preparation  of  alcohol.  Alcohol  is  prepared  com- 
mercially from  starch  obtained  from  corn  or  potatoes.  The  starch 
is  first  converted  into  a  sugar  known  as  maltose,  by  the  action 
of  malt,  a  substance 
prepared  by  moisten- 
ing barley  with  water, 
allowing  it  to  sprout, 
and  then  drying  it. 
This  sugar,  like  dex- 
trose, breaks  down  into 
alcohol  and  carbon  di- 
oxide in  the  presence 
of  yeast.  The  resulting 
alcohol  is  separated  by 
fractional  distillation.  FIG.  115.  Laboratory  preparation  of  alcohol 

2.  Properties.  Ethyl  alcohol  is  a  colorless  liquid  with  a 
pleasant  odor,  and  is  an  excellent  solvent  for  many  organic 
substances.  It  boils  at  78.3°.  It  is  sometimes  used  as  a 


230  FIRST  COURSE  IN  CHEMISTEY 

fuel,  since  its  flame  is  very  hot  and  does  not  deposit  car- 
bon, as  the  flame  from  oil  does.  When  taken  into  the  sys- 
tem in  small  quantities,  it  causes  intoxication ;  in  larger 
quantities  it  acts  as  a  poison.  The  ordinary  alcohol  of  the 
druggist  contains  about  95  per  cent  alcohol  and  5  per  cent 
water.  A  solution  containing  99  per  cent  or  more  of  alcohol 
is  called  absolute  alcohol.  When  alcohol  is  heated  with  sul- 
fur ic  acid,  a  low-boiling  inflammable  liquid  known  as  ether 
is  formed:  2  C>H(OH -- »- (C^.O  +  H,O 

This  is  largely  used  as  an  anaesthetic  in  surgical  operations. 
Denatured  alcohol.  Ordinary  alcohol  (95  per  cent)  sells 
at  about  $2.60  per  gallon.  Of  this,  $2.10  is  a  tax  imposed 
by  the  government.  By  an  act  of  Congress  in  1906  the 
tax  was  removed  from  denatured  alcohol;  that  is,  alcohol 
mixed  with  some  substance  which  renders  it  unfit  for  use 
as  a  beverage  but  does  not  impair  its  use  for  manufacturing 
purposes.  The  substances  ordinarily  used  for  this  purpose 
are  methyl  alcohol,  benzine,  and  a  compound  prepared  by 
heating  bones  and  known  as  pyridine. 

Alcoholic  liquors.  All  alcoholic  liquors  are  made  by  alcoholic 
fermentation.  Wine  is  made  by  the  fermentation  of  the  dex- 
trose in  grape  juice  and  contains  from  5  to  15  per  cent  by 
volume  of  alcohol.  Beer  is  made  from  maltose  formed  by  the 
action  of  malt  upon  starch  obtained  from  various  grains,  chiefly 
barley.  It  contains  from  3  to  5  per  cent  by  volume  of  alcohol. 
Whisky  contains  about  50  per  cent  by  volume  of  alcohol  and 
is  made  from  starch  by  a  process  very  similar  to  that  described 
under  the  commercial  preparation  of  alcohol.  Almost  any  sac- 
charine liquid,  such  as  cider  and  the  juices  of  fruits  in  general, 
undergoes  alcoholic  fermentation  when  exposed  to  air. 

Alcoholic  liquors,  as  well  as  pure  alcohol,  are  taxed  by  the 
government.  In  the  year  1912-1913  the  revenue  collected  from 
this  source  amounted  to  over  $223,000,000. 


ALCOHOLS  231 

Chemical  changes  in  bread-making.  The  average  com- 
position of  wheat  flour  is  as  follows: 

Water 13.8% 

Protein  (nitrogenous  matter) 7.9% 

Fats 1-4% 

Starch 76.4% 

Mineral  matter 0.5% 

In  making  bread,  flour  is  mixed  with  water,  yeast, 
and  a  little  sugar,  and  the  resulting  dough  is  set  aside  hi 
a  warm  place  for  a  few  hours.  The  yeast  first  causes  the 
sugar  to  undergo  alcoholic  fermentation.  The  carbon  diox- 
ide formed  escapes  through  the  dough,  making  it  light  and 
porous.  The  yeast  plant  thrives  best  at  about  30° ;  hence 
the  necessity  for  keeping  the  dough  in  a  warm  place.  In 
baking  bread  the  heat  expels  the  alcohol  and  also  expands 
the  bubbles  of  carbon  dioxide  caught  in  the  dough,  causing 
it  to  become  porous  and  making  the  bread  light. 

Preservatives.  We  have  observed  that  the  changes  taking 
place  in  the  souring  of  milk  and  the  changing  of  sugar  into 
alcohol  are  caused  by  low  forms  of  organisms,  the  cells  of 
which  are  present  in  the  air.  Many  other  similar  changes, 
such  as  putrefaction,  are  due  to  the  same  causes.  All  these 
changes  may  be  prevented  in  one  of  the  following  ways : 

1.  By  keeping  the  substance  at  such  a  low  temperature 
that  the  organism  causing  the  change  cannot  thrive  (cold- 
storage). 

2.  The   substance   may  be  heated  so  as  to  destroy  all 
organisms  present,  and  then  sealed  air-tight  in  a  suitable 
container.    This  is  the  method  used  in  canning  vegetables 
and  in  preserving  grape  juice  and  condensed  milk. 

3.  Some  substance  may  be  added  which  in  small  amounts 
will  destroy  the  organisms  causing  the  change  or  will  prevent 
their  growth.    Such  a  substance  is  known  as  a  preservative. 


232  FIRST  COURSE  IN  CHEMISTRY 

Whether  or  not  preservatives  should  be  permitted  in  foods 
is  a  much  debated  question.  Some  people  maintain  that  any 
substance  which  is  powerful  enough  to  prevent  the  growth 
of  the  organisms  must  have  an  injurious  action  upon  diges- 
tion. The  federal  government  at  present  allows  the  use 
of  sodium  benzoate  (a  white  solid  made  from  a  hydrocar- 
bon present  in  coal  tar)  in  such  foods  as  jellies,  jams,  and 
catchup,  which  are  not  consumed  immediately  upon  the 
opening  of  the  container.  If  this  preservative  is  used, 
however,  the  labels  on  the  containers  must  state  the 
amount  present. 

Some  derivatives  of  coal  tar.  In  discussing  the  manufac- 
ture of  coal  gas,  it  was  stated  that  from  the  coal  tar  formed 
in  the  process  there  is  obtained  a  large  number  of  impor- 
tant compounds.  These  are  often  spoken  of  collectively  as 
the  coal-tar  compounds.  It  is  possible  here  to  mention  only 
a  few  of  these. 

(1)  Benzene  (C6H6)  and  (2)  toluene  (C7Hg)  are  highly  inflam- 
mable, colorless  liquids ;  (3)  naphthalene  (C10Hg)  and  (4)  an- 
thracene (C14H10)  are  white,  solid  hydrocarbons,  which  are  used 
in  the  preparation  of  the  two  dyes,  indigo  and  alizarin.  These 
dyes  were  formerly  obtained  from  vegetable  sources,  but  are 
now  manufactured  at  low  cost.  Ordinary  moth  balls  are  nearly 
pure  naphthalene.  (5)  Phenol,  or  carbolic  acid  (C6H5OH),  is  a 
"white,  crystalline  solid,  very  caustic  and  poisonous.  (6)  Cresol 
(C7H7OH)  is  a  good  disinfectant,  and  is  the  basis  of  most  of 
the  common  disinfectants  and  sheep  dips  now  on  the  market. 

Each  of  the  above  compounds  serves  as  the  starting  ma- 
terial from  which  many  other  useful  compounds  are  prepared. 
Thus,  benzene  when  treated  with  nitric  acid  gives  nitrobenzene 
(C6H5N02),  and  this  on  reduction  yields  aniline  (C6H5NH2). 
Aniline  is  a  nearly  colorless  liquid,  and  from  it  are  prepared  a 
large  number  of  dyes  of  all  shades  and  colors,  known  as  the 
aniline  dyes.  Toluene  when  oxidized  forms  benzoic  acid,  the 


COAL-TAR  COMPOUNDS  233 

sodium  salt  of  which  (sodium  benzoate)  is  used  as  a  food  pre- 
servative. When  phenol  is  heated  with  formaldehyde  there  are 
obtained  products  known  commercially  as  bakelite  and  condens- 
ite.  These  are  useful  materials  for  making  buttons,  umbrella 
handles,  pipestems,  and  insulators  in  electrical  apparatus. 

Coal-tar  compounds  in  foods.  Much  discussion  has  arisen 
in  regard  to  the  use  of  coal-tar  compounds  in  foods.  The 
federal  government  has  selected  seven  different  aniline 
dyes  of  different  colors  the  use  of  which  is  permitted  in 
such  foods  as  candies  and  butter.  As  already  stated 
(p.  232),  the  use  of  sodium  benzoate  as  a  preservative  is 
allowed  under  certain  restrictions.  Saccharine,  a  white  solid 
prepared  from  toluene  and  500  times  as  sweet  as  sugar,  was 
formerly  permitted  in  foods,  but  in  1912  the  government 
forbade  the  further  use  of  it.  Vanillin,  identical  with  the 
compound  prepared  from  vanilla  beans,  and  coumarin,  which 
has  an  odor  similar  to  vanillin,  are  both  used,  in  artificial 
vanilla  extracts,  but  when  they  are  so  used,  the  label  on  the 
container  must  state  the  fact.  It  is  well  to  keep  in  mind  that 
all  such  substances  are  not  foods  and  are  used  for  purposes 
other  than  nutrition. 

EXERCISES 

1.  What  is  the  meaning  of  the  term  carbohydrate  (see  dictionary)? 

2.  Can  you  tell  the  difference  between  pure  sugar  obtained -from 
sugar  cane  and  that  obtained  from  the  sugar  beet? 

3.  It  is   often    said   that   milk    sours    readily  during   thunder 
showers.    What  would  you  say  as  to  the  truth  of  this  statement? 

4.  Why  do  we  use  corn  rather  than  dextrose  in  making  alcohol? 

5.  In  separating  alcohol  from  water  by  distillation,  which  dis- 
tills over  first  ? 

6.  Why  does  the  government  permit  the  use  of  a  preservative 
(sodium  benzoate)  in  catchup  but  not  in  milk  ? 

7.  What  weight  of  starch  is  necessary  in  making  1  ton  of  pure 
dextrose?    Ans.  1799.93  Ib. 


234  FIRST  COURSE  IN  CHEMISTRY 

8.  How  could  you  tell  the  difference  between  methyl  and  ethyl 
alcohols  ? 

9.  Yeast  is  often  added  in  preparing  household  beverages.   Why 
is  it  added?    What  substance  will  be  present  in  the  beverage  so 
prepared  ? 

10.  Why  is  sugar  (or  molasses)  added  in  making  bread? 

11.  Alcohol  and  gasoline  boil  at  about  the  same  temperature  and 
both  are  combustible.     Why  not  use  alcohol  as  a  fuel  in  place  of 
gasoline  ? 

12.  Can  you  suggest  a  method  for  obtaining  alcohol  from  wood? 

TOPICS  FOR  THEMES 

Glucose  (Lassar-Cohn,  Chemistry  in  Daily  Life). 

The  refining  of  sugar  (Wiley,  Foods  and  their  Adulteration). 

Alcohol  in  the  industries  (Duncan,  Chemistry  of  Commerce). 

Uses  of  cellulose  (Duncan,  Chemistry  of  Commerce ;  Bird,  Modern 
Science  Reader). 

Paper-making  (Rogers  and  Aubert,  Industrial  Chemistry;  Lassar- 
Cohn,  Chemistry  in  Daily  Life). 


CHAPTER  XXIV 
ORGANIC  ACIDS  ;   FATS ;   OILS  ;   PROTEINS 

Organic  acids.  A  great  number  of  acids  are  known  which 
are  composed  of  carbon,  oxygen,  and  hydrogen,  and  as  a 
group  these  are  called  organic  acids.  Like  the  hydrocarbons, 
they  can  be  arranged  in  series,  one  of  the  most  important  of 
which  is  known  as  the  fatty-acid  series.  A  few  of  the  most 
important  acids  of  this  series  are  given  in  the  following  table. 
They  are  all  monobasic  —  a  fact  indicated  in  the  formula  by 
separating  the  replaceable  hydrogen  atom  from  the  rest  of 
the  molecule. 


SOME  FATTY  ACIDS 

II  •  CH<X formic  acid,  a  liquid  boiling  at  100° 

V 


II  •  C0H3(X acetic  acid,  a  liquid  boiling  at  118° 


II  •  C4H7O2 butyric  acid,  a  liquid  boiling  at  163° 

II  •  C16H31O0 palmitic  acid,  a  solid  melting  at  62° 

II  •  C18H35O2 stearic  acid,  a  solid  melting  at  69° 

H'CnH2n_iO2 general  formula 

Of  these,  acetic  acid  deserves  special  mention. 

Acetic  acid  (H  •  C2H302).  This  is  the  acid  which  imparts 
the  sour  taste  to  vinegar.  It  is  prepared  commercially  by 
the  distillation  of  wood  (p.  196).  It  is  a  colorless  liquid 
and  has  a  strong,  pungent  odor.  When  anhydrous,  it 
crystallizes  as  a  white  solid  which  melts  at  18°  and  closely 
resembles  ice  in  appearance ;  hence  the  name  glacial  acetic 
acid.  Many  of  the  salts  of  acetic  acid  are  well-known  com- 
pounds. Thus,  lead  acetate  (Pb(C2H8O2)a  •  3  H2O)  is  the 
white  solid  known  as  sugar  of  lead. 

235 


236  FIRST  COURSE  IN  CHEMISTRY 

Vinegar.  As  is  well  known,  when  cider  is  exposed  to  the  air 
it  is  gradually  transformed  into  vinegar.  Two  changes  are 
involved  in  the  process  :  (1)  the  sugar  in  the  cider  first  under- 
goes alcoholic  fermentation,  forming  hard  cider,  which  contains 
from  4  to  8  per  cent  of  alcohol ;  (2)  the  alcohol  is  then  oxidized 
to  acetic  acid,  the  necessary  oxygen  coming  from  the  air.  This 
oxidation  is  brought  about  through  the  action  of  the  micro- 
organism known  as  Mycoderma 
aceti.  This  organism  is  present  in 
the  so-called  mother  of  vinegar. 
The  oxidation  of  alcohol  into  acetic 
acid  through  the  agency  of  the 
Mycoderma  aceti  is  known  as  acetic 
fermentation,  and  may  be  repre- 
sented as  follows : 

C2H5OH  +  02  — >  H  •  C2H302  +  H20 

The  manufacture  of  vinegar.    The 

old  method  of  making  vinegar  con- 
sisted simply  in  storing  cider  in 
barrels  until  the  fermentation  was 
complete.  In  the  modern  method 
the  change  is  brought  about  in  a 
few  hours,  a  large  cask,  known  as 
a  generator,  being  used  (Fig.  116). 
This  is  filled  loosely  with  beech 
FIG.  116.  A  vinegar  generator  shavings.  Vinegar  is  first  sprayed 

into  the  top  of  the  cask  in  order  to 

introduce  the  Mycoderma  aceti.  The  organism  attaches  itself 
to  the  wood  shavings,  which  are  used  because  they  present  a 
large  surface.  Next  a  dilute  solution  of  alcohol  (hard  cider,  in 
the  case  of  cider  vinegar)  is  sprayed  into  the  top  of  the  cask 
while  air  is  admitted  at  the  bottom  A,  A.  In  this  way  the 
alcohol  and  oxygen  are  brought  into  intimate  contact,  and  the 
oxidation  takes  place  rapidly  as  the  liquid  trickles  down  over 
the  shavings.  The  resulting  vinegar  is  drawn  off  at  the  bottom,  B, 
of  the  cask.  Instead  of  starting  with  cider,  one  may  use  almost 


OKGANIC  ACIDS;  FATS;  OILS;  PBOTEINS    237 

any  substance  which  contains  starch  or  sugar,  these  compounds 
first  being  changed  into  alcohol,  as  explained  in  the  manufacture 
of  alcohol.  In  this  way  are  prepared  malt  vinegar  from  starch 
and  sugar  vinegar  from  sugar  residues.  The  cheapest  vinegar 
is  made  from  pure  dilute  alcohol,  and  is  known  as  distilled 
vinegar.  It  is  colorless  and  leaves  no  residue  upon  evaporation. 
The  federal  law  requires  that  all  vinegar  shall  contain  not 
less  than  4  per  cent  acetic  acid.  In  addition  to  the  acid,  vinegar 
prepared  from  fruits  and  grains  contains  certain  solids  derived 
from  the  source  materials.  It  is  by  studying  the  character  of 
these  solids  left  upon  evaporating  a  sample  of  vinegar  that 
the  chemist  is  able  to  determine  the  source  of  the  vinegar. 

Acids  belonging  to  other  series.  In  addition  to  the  fatty 
acids,  the  following  deserve  special  mention. 

Tartaric  acid  (Ha-C4H406).  This  is  a  white  solid  and 
occurs  in  many  fruits  either  in  the  free  state  or  in  the  form 
of  its  salts.  The  acid  potassium  salt  (KHC4H4O6)  occurs  in 
the  juice  of  grapes.  When  the  juice  ferments  in  the  manu- 
facture of  wine,  this  salt,  being  insoluble  in  alcohol,  separates 
on  the  sides  of  the  cask  and  in  this  form  is  known  as  argol. 
When  purified,  it  forms  a  white  solid,  which  is  sold  under 
the  name  of  cream  of  tartar  and  is  used  in  baking  powders 
(p.  330).  The  acid  itself  is  often  used  hi  soft  drinks  (pop). 

Citric  acid  (H3  •  C6H507).  This  acid  occurs  in  many  fruits, 
especially  in  lemons.  It  is  a  white  solid,  soluble  in  water,  and 
is  often  used  as  a  substitute  for  lemons  in  making  lemonade. 

Oleic  acid  (H  •  C18H3302).  The  derivatives  of  this  acid 
constitute  the  principal  part  of  many  oils  and  liquid  fats. 
The  acid  itself  is  an  oily  liquid. 

Fats  and  oils.  The  hydrogen  of  acids  can  be  replaced 
not  only  by  metals  but  by  hydrocarbon  radicals  as  well. 
The  resulting  compounds  are  termed  esters.  The  main  con- 
stituents of  the  common  fats  and  oils,  such  as  butter,  lard, 
and  olive  oil,  are  esters  of  oleic,  palmitic,  and  stearic  acids 


238  FIKST  COUESE  IN  CHEMISTRY 

and  are  known  respectively  as  olein,  palmitin,  and  stearin. 
The  radical  present  in  these  esters  is  C3H5.  It  is  trivalent 
and  is  known  as  the  glyceryl  radical,  since  it  is  present  in 
glycerin  (C3H5(OH)3).  Since  the  glyceryl  radical  is  triva- 
lent, and  since  oleic,  palmitic,  and  stearic  acids  are  all  mono- 
basic, it  is  evident  that  three  molecules  of  each  acid  must 
enter  into  the  formation  of  each  molecule  of  the  ester  derived 
from  it.  The  relation  in  composition  between  these  acids  and 
the  corresponding  esters  is  shown  in  the  following  formulas: 

II  •  C18H3302  (oleic  acid) C8H6(C18H88O2)8  (olein) 

II .  C16H31O2  (palmitic  acid)  .  .  .  C3H5(C1CH31O2)3  (palmitin) 
H  .  C18H35O2  (stearic  acid)  ....  C3HC(C18H3-O2)8  (stearin) 

Olein  is  a  liquid  and  is  the  main  constituent  of  oils  such 
as  olive  oil.  Palmitin  and  stearin  are  white  solids.  Beef 
suet  is  principally  stearin. 

Butter  fat  and  oleomargarine.  While  butter  fat,  like 
other  fats,  consists  principally  of  olein,  palmitin,  and  stearin, 
its  characteristic  flavor  is  due  to  the  presence  of  a  small 
amount  (about  8  per  cent)  of  the  fat  butyrin,  which  is  an 
ester  of  butyric  acid  and  has  the  formula  CfH6(C4H7O2)8. 
Oleomargarine  differs  from  butter  mainly  in  the  fact  that  a 
smaller  amount  of  butyrin  is  present.  It  is  made  from  the 
fats  obtained  from  cattle  and  hogs.  Sometimes  cottonseed 
oil  is  also  added.  These  fats  are  churned  with  milk,  or 
mixed  with  a  small  amount  of  butter,  in  order  to  furnish 
sufficient  butyrin  to  give  the  butter  flavor. 

In  appearance  oleomargarine  differs  from  most  butter  in 
being  nearly  colorless.  While  it  is  a  common  practice  to  color 
butter  artificially,  the  federal  law  permits  the  coloring  of  oleo- 
margarine only  upon  the  payment  of  a  tax  of  10  cents  for 
each  pound  colored.  Many  of  the  states,  however,  have  laws 
forbidding  the  sale  of  oleomargarine  that  is  artificially  colored, 
even  though  the  federal  tax  has  been  paid. 


OKGANIC   ACIDS;   FATS;   OILS;  PKOTEINS    239 

Changing  oils  into  solid  fats.  It  will  be  noted  that  stearin 
differs  from  olein  in  composition  in  that  it  contains  6  atoms 
of  hydrogen  more  in  each  molecule.  Now  if  hydrogen  is 
brought  in  contact  with  olein  under  proper  conditions  and 
in  the  presence  of  a  suitable  catalytic  agent  (finely  divided 
nickel  is  used),  the  olein  takes  up  the  additional  hydrogen 
and  is  changed  into  the  solid  stearin.  It  is  possible  in  this 
way  to  change  the  oils  into  solid  fats.  Certain  commercial 
fats  used  in  cooking,  such  as  that  sold  under  the  name  of 
Crisco,  are  made  by  this  process  from  the  comparatively 
inexpensive  cottonseed  oil. 

The  proteins.  The  term  protein  is  applied  to  a  large  class 
of  complex  nitrogenous  compounds  which  are  everywhere 
abundant  in  animal  and  vegetable  organisms  and  which 
constitute  the  principal  part  of  the  tissues  of  the  living  cell. 
The  casein  of  milk,  gluten  of  flour,  and  albumin  of  egg  will 
serve  as  examples  of  typical  protein  matter.  The  proteins 
all  contain  nitrogen,  carbon,  hydrogen,  and  oxygen,  and 
some  contain  sulfur  and  phosphorus  in  addition. 

EXERCISES 

1.  Since   potassium  'bitartrate  (cream    of   tartar)  is  soluble  in 
water,  why  does  it  crystallize  out  when  grape  juice  ferments? 

2.  How  could  you  prevent  cider  from  changing  into  vinegar? 

3.  What  is  hard  cider?    How  does  it  differ  in  composition  from 
sweet  cider  ? 

4.  Supposing  that  95  per  cent  of  the  sugar  undergoes  alcoholic 
fermentation,  what  weight  of  alcohol  can  be  prepared  from  100  Ib.  of 
dextrose?  Ans.  48.6  Ib.    What  weight  of  acetic  acid  can  be  prepared 
from  this  alcohol?    Ans.  63.4  Ib. 

TOPICS  FOR  THEMES 

Catalysis  (Duncan,  Chemistry  of  Commerce). 

Vinegar-making.  (Write  to  United  States  Department  of  Agricul- 
ture, Washington,  D.  C.,  for  bulletins.) 


CHAPTER  XXV 
FOODS 

Composition  of  foods.  While  the  compounds  present  in 
our  foods  are  very  numerous  and  often  exceedingly  com- 
plex, yet  they  may  all  be  included  under  a  few  general 
heads,  namely,  proteins,  fats,  carbohydrates,  mineral  matter, 
and  water.  Since  the  mineral  matter  is  left  as  a  residue 
when  the  food  is  burned,  it  is  listed  as  ash  in  reporting  the 
analysis  of  foods.  The  composition  of  the  more  common 
foods  is  given  in  the  table  on  the  opposite  page. 

Function  of  foods.    Foods  have  a  twofold  function 

1.  They  provide  the  material  for  the  growth  of  the  body 
as  well  as  for  the  repair  of  worn-out  tissues. 

2.  They  furnish  the  necessary  energy  for  muscular  work 
and  for  maintaining  the  heat  of  the  body. 

Broadly  speaking,  it  may  be  said  that  the  first  of  these 
functions  is  performed  by  the  protein  matter  of  our  foods 
together  with  certain  mineral  salts,  while  the  carbohydrates 
and  fats  and,  to  a  certain  extent  the  proteins  also,  are  en- 
ergy producers.  The  mineral  matter  supplies  the  material 
for  building  up  the  solid  tissues  of  the  body,  such  as  the 
bones,  and  has  in  addition  more  complex  functions.  Water 
serves  to  assist  in  promoting  chemical  action  and  to  carry 
material  from  one  part  of  the  body  to  another. 

While  the  different  classes  of  food  materials  are  to  a  cer- 
tain extent  interchangeable,  experiments  show  that  a  proper 
mixture  of  these  materials  is  essential  to  health.  Of  course, 

240 


FOODS 


241 


AVERAGE  COMPOSITION  OF  EDIBLE  PORTION  OF  TYPICAL  FOODS 
EXPRESSED  IN  GRAMS  PER  100  GRAMS  OF  FOOD 


FOOD 

WATER 

PROTEIN 

FAT 

CARBO- 
HYDRATE 

ASH 

FUEL 
VALUE 
(Cal.  per 
100  g.) 

Almonds  

48 

21.0 

549 

173 

20 

647 

Apples 

84  G 

0  4 

0  5 

14  2 

03 

63 

Asparagus  .... 
Bacon  (smoked)  .  . 
Bananas  

94.0 
20.2 
75.3 

1.8 
9.9 
1.3 

0.2 
64.8 
0  6 

3.3 

22  0 

0.7 
5.1 

08 

22 
623 
99 

Beans  (dried)  .  .  . 
Beans  (string)  .  .  . 
Beef  (lean  steak)  .  . 
Beef  (slightly  fat)  . 
Beets  .  .  . 

12.6 

89.2 
70.0 
73.8 
87  5 

22.5 
2.3 
21.0 
22.1 
1  6 

1.8 
0.3 
7.9 
2.9 
0  1 

59.6 
7.4 

9  7 

3.5 
0.8 
.1 

1 

345 
42 
155 
115 
46 

Bread  (corn)  .  .  . 
Bread  (graham)  .  . 
Bread  (white)  .  .  . 
Butter  
Cabbage  . 

38.9 
35.7 
35.3 
11.0 

91  5 

7.9 
8.9 
9.2 
1.0 
1  6 

4.7 
1.8 
1.3 
85.0 
03 

46.3 
52.1 
53.1 

5  6 

2.2 
.5 
.1 
3.0 
1  0 

259 
260 
260 
769 
32 

Carrots  
Celery  .... 

88.2 
94  5 

1.1 
1  1 

0.4 
0  1 

9.3 
33 

1.0 
1  0 

45 
19 

Chestnuts  .... 
Chicken  
Codfish  (fresh)  .  .  . 
Corn  (green)  .  .  . 
Dates  . 

45.0 
63.7 
82.6 
75.4 
13  8 

6.2 
19.3 
15.8 
3.1 
1  9 

5.4 
16.3 
0.4 
1.1 

2  5 

42.1 

19.7 
70  6 

1.3 
1.0 
1.2 
0.7 
1  2 

242 
224 
67 
101 
313 

Eggs 

73  7 

148 

105 

10 

154 

Figs  .  .  . 

18  8 

4  3 

0  3 

74  2 

2  4 

317 

Ham  (lean,  smoked)  . 
Lettuce 

53.5 

94  7 

20.2 
1  2 

20.8 
0  3 

2  0 

5.5 
09 

268 
16 

Macaroni  . 

784 

30 

1  5 

15  8 

1  3 

89 

Milk  
Oatmeal 

87.0 
7  3 

3.3 
16  1 

4.0 

7  2 

5.0 

67  5 

0.7 
1  9 

69 

400 

Olive  oil  .  . 

100  0 

900 

Oranges 

869 

0  8 

0  2 

11  6 

05 

51 

Peaches 

89  4 

0  7 

0  1 

9  4 

04 

41 

Peanuts  

9.2 

25  8 

38  6 

244 

2.0 

548 

Peas  (green)  .... 
I'luius  
Potatoes  .... 

74.6 

78.4 
783 

7.0 
1.0 
2  2 

0.5 
0  1 

16.9 
20.1 
18  4 

1.0 
0.5 
1  0 

100 

84 
83 

Prunes  (dried)  . 
Raisins      
Rice  .... 

22.3 
14.6 
12  3 

2.1 
2.6 

8  0 

3.3 
03 

73.3 
76.1 
79  o 

2.3 
3.4 
04 

302 
345 
351 

Salmon  

64  6 

21  2 

J9  g 

1  4 

200 

Spinach 

92  3 

2  1 

0  3 

3  2 

2  1 

24 

Strawberries  .  .  . 
Tomatoes  
Turnips  

90.4 
94.3 
89.6 

1.0 
0.9 
1  3 

0.6 
0.4 
0  2 

7.4 

3.9 
8  1 

0.6 
0.5 
0.8 

39 
23 
40 

Wheat  flour  .... 

11.9 

13.3 

1.5 

72.7 

0.6 

357 

These  values  are  taken  from  Bulletin  Xo.  28,  office  of  Experiment  Station, Washing- 
ton, D.  C.  The  fuel  values  are  obtained  from  the  following  formula  : 

Cal.  in  100g.  =  4/)  +  9F+4C,  in  which  P,  F,  and  C  represent  respectively  the 
number  of  grams  of  protein,  fat,  and  carbohydrates  in  100  g.  of  the  food  (p.  242). 


242  FIRST  COURSE  IN  CHEMISTRY 

it  is  true  that  one  can  live  for  many  clays  on  a  purely 
protein  diet,  or  on  a  diet  purely  of  fats  and  carbohydrates ; 
in  fact,  persons  have  been  known  to  live  for  many  days 
without  any  food  whatever  (other  than  water).  In  all  such 
cases  the  body  derives  the  necessary  materials  from  the 
surplus  supply  always  stored  up  in  the  normal  body. 

The  energy  value  of  foods.  Experiments  show  that  the  heat 
of  the  body,  as  well  as  the  energy  used  in  muscular  work, 
results  from  the  oxidation  of  food  materials.  The  foods, 
when  eaten,  undergo  complex  changes  in  which  the  insoluble 
portions  are  converted  into  soluble  compounds.  These  are 
absorbed  into  the  system  and  then  either  undergo  oxidation 
directly  or  are  temporarily  built  into  tissues  which  later  un- 
dergo oxidation.  In  this  process  most  of  the  carbon  is  finally 
changed  into  carbon  dioxide  and  exhaled  from  the  lungs, 
while  the  hydrogen  is  changed  into  water.  The  nitrogen  is 
excreted  largely  in  the  form  of  urea  (CO(NH2).i). 

Broadly  speaking,  foods  may  be  regarded  as  fuel  from  the 
oxidation  of  which  in  the  body  the  energy  necessary  for  the 
bodily  requirements  is  set  free.  In  the  study  of  foods  it  is  con- 
venient, therefore,  to  use  their  fuel  values  (heats  of  combus- 
tion) as  a  basis  of  comparison.  These  values  are  determined 
in  the  calorimeter  and  expressed  in  large  calories  (Cal.), 
which  are  1000  times  as  large  as  the  small  calorie  (cal.). 

Now  experiments  show  that  the  body  obtains  from  each 
of  the  three  classes  of  foods,  when  absorbed  in  the  system 
and  oxidized,  approximately  the  following  fuel  values : 


CLASS  OF  FOODS 

CALORIES 

I'KK    (Jl.'AM 

C  A  I.OKIKS 
PER   POUND 

Carbohydrates 

4 

1815 

Fats    

9 

4082 

Proteins        

4 

1815 

FOODS 


243 


The  amount  and  nature  of  foods  necessary  for  health. 
Many  studies  have  been  made  in  order  to  find  out  just  how 
much  and  what  kind  of  food  is  best  adapted  for  the  pres- 
ervation of  health.  Evidently  many  conditions,  such  as 
one's  age,  weight,  and  occupation,  and  the  climate  in  which 
one  lives,  enter  into  the  problem. 

Since  the  fats  and  carbohydrates  have  nearly  the  same 
function,  it  is  sufficient  in  stating  food  requirements  for  a 
period  of,  say,  twenty-four  hours,  to  give  simply  the  weight 
of  protein  necessary,  together  with  the  total  fuel  value. 
The  difference  between  the  total  fuel  value  and  that  of 
the  required  protein  gives  the  number  of  calories  to  be 
supplied  from  fats  and  carbohydrates.  Such  a  mixture  of 
these  two  food  materials  is  selected  as  will  suit  the  taste, 
provided  only  that  the  fuel  value  of  the  mixture  together 
with  the  fuel  value  of  the  protein  equals  the  total  fuel 
value  required.  The  following  dietary  standards  proposed 
by  Atwater  are  generally  accepted.  They  give  the  food 
requirements  for  a  period  of  twenty-four  hours. 


CHARACTER  OF  INDIVIDUAL 

P  ROT  KINS 
REQUIRED 

FUEL  VALUE 
REQUIRED 

Man  with  very  hard  muscular  work  (wood- 
chopper  footl)all  player^ 

175  o'. 

5500  Cal. 

Man  with  moderately  active  muscular  work 
Man  with  light  to  moderate  muscular  work 
Man  at  sedentary,  or  woman  with  moder- 
ately active  work    

125  g. 
112  g. 

100  g. 

3400  Cal. 
3050  Cal. 

2700  Cal. 

Man  at  rest  or  woman  with  light  muscular 

90  g. 

2450  Cal. 

Boy  15  to  16  years       

108  g. 

3060  Cal. 

Boy  13  to  14  years  or  girl  15  to  16   .     .     . 
Boy  12  to  13  years  or  girl  14  to  15    . 
Boy  10  to  11  years  or  girl  10  to  12    .     .     . 
Boy  6  to  9  years      

100  g. 
87  g. 
75  g. 
62  g. 

2720  Cal. 
2380  Cal. 
2040  Cal. 
1700  Cal. 

244 


FIRST   COURSE  IN  CHEMISTRY 


The  calculation  of  diets.  With  the  aid  of  the  standards 
given  above  and  the  table  giving  the  composition  of  foods 
(p.  241)  it  is  possible  to  select  suitable  diets  in  the  case  of 
any  individual. 

The  laws  of  conservation  of  energy  and  matter  in  the 
human  body.  In  studying  the  food  requirements  of  the 
human  body,  many  experiments  have  been  made  to  find  out 


FIG.  117.    Exterior  of  a  respiration  calorimeter 

whether  or  not  the  laws  holding  good  in  the  inanimate 
world  (namely,  the  laws  of  conservation  of  matter  and 
energy)  hold  good  also  in  the  processes  taking  place  in 
the  living  organism.  So  far  as  the  law  of  conservation  of 
matter  is  concerned,  the  question  can  be  answered  by  care- 
fully comparing  the  air  breathed  and  food  eaten  (the  intake*) 
with  the  products  stored  up  in  the  body  and  with  the  prod- 
ucts exhaled  or  excreted  (the  output). 


FOODS  245 

In  order  to  measure  the  energy  changes,  however,  one 
must  be  able  to  determine  the  amount  of  heat  evolved. 
This  is  done  by  means  of  the  respiration  calorimeter.  This 
is  a  chamber  large  enough  to  enable  one  to  live  in  it,  and 
constructed  with  double  walls  of  nonconducting  material 
so  as  to  prevent  loss  of  heat  through  radiation  —  in  fact, 
it  is  a  large  calorimeter  (Fig.  36).  Provision  is  made  for 
introducing  air  and  food. 

With  the  necessary  precautions  it  is  possible  in  this  way 
to  measure  the  heat  generated  in  the  body  of  a  person  liv- 
ing within  the  calorimeter.  The  experiments  show  that  the 
changes  in  matter  and  energy  which  take  place  in  the 
human  body  are  in  accord  with  the  laws  of  the  conservation 
of  matter  and  energy. 

Fig.  117  shows  the  exterior  of  a  respiration  calorimeter;  also 
the  apparatus  for  analyzing  the  air  admitted,  and  measuring 
the  heat  evolved  by  the  person  in  the  calorimeter. 

The  cost  of  food  as  related  to  its  nutritive  value.  It  will 
be  noted  that  the  food  requirements  for  the  body  are  stated 
simply  in  terms  of  protein,  carbohydrates,  and  fats.  It  is 
very  evident  that  the  cost  of  say  100  g.  of  protein  will  vary 
according  to  the  source  of  the  protein.  For  example,  pro- 
tein obtained  in  the  form  of  tenderloin  steak  will  cost  over 
8  times  as  much  as  an  equal  weight  of  protein  obtained 
from  dried  beans.  Again,  there  is  as  much  nutriment  in  1  Ib. 
of  wheat  flour  (cost,  about  4  cents)  as  in  3|  qt.  of  oysters 
(cost,  about  $1.40).  It  is  interesting  that  4  cents'  worth  of 
corn  meal  or  oatmeal  contains  sufficient  protein  and  total 
fuel  value  to  supply  the  needs  of  the  body  for  twenty-four 
hours.  A  diet  of  this  kind,  however,  is  not  agreeable,  since 
we  demand  food  that  is  more  highly  palatable  and  more 
attractive  to  the  senses.  By  finding  the  selling  prices  of 


246  FIKST  COUKSE  IK  CHBMISTEY 

the  various  foods  given  in  the  table  on  page  241,  one  can 
easily  determine  the  relative  cost  of  various  food  materials 
obtained  from  different  sources. 

Animal  nutrition  as  contrasted  with  plant  nutrition.  Plants 
build  up  their  tissues  from  water,  nitrogenous  matter,  and 
mineral  salts  absorbed  from  the  soil,  and  carbon  dioxide 
absorbed  from  the  air.  In  this  process  a  large  amount  of 
energy  derived  from  the  sun's  rays  is  absorbed. 

Animals,  on  the  other  hand,  live  upon  these  complex 
vegetable  substances,  and  in  the  course  of  their  assimilation 
the  processes  which  took  place  in  their  formation  in  the  plant 
are  largely  reversed.  The  carbon  and  hydrogen  are  again 
liberated,  principally  as  carbon  dioxide  and  water,  the  nitro- 
gen is  secreted  largely  in  the  form  of  urea  CO(NH2)2,  while 
the  energy  stored  up  in  the  plant  is  utilized  as  a  source  of 
heat  and  for  muscular  work.  In  general,  it  may  be  said, 
therefore,  that  the  processes  involved  in  plant  and  animal 
nutrition  are  opposite  in  character :  the  plant  builds  up  com- 
plex tissues  out  of  simple  compounds  with  energy  absorption  ; 
the  animal  decomposes  these  tissues  into  simple  compounds  with 
evolution  of  energy. 

EXERCISES 

1.  Is  the  ash  obtained  in  burning  a  food  present  in  that  form 
in  the  food  ? 

2.  Why  do  different  nations  use  different  classes  of  foods  ? 

3.  Mention  some  foods  that  are  rich  in  water;   in  protein;   in 
fats;   in  carbohydrates. 

4.  Ascertain  the  current  prices  of  common  foods,  then  work  out 
a  list  of  the  most  economic  ones  to  use  as  a  source  of  protein ;  of 
carbohydrates  ;   of  fats. 

TOPICS  FOR  THEMES 

Butter  and  oleomargarine  (Wiley,  Foods  and  their  Adulteration). 
The  cottonseed  oil  industry  (Wiley, Foods  and  their  Adulteration). 


CHAPTER  XXVI 
THE  PHOSPHORUS  FAMILY 


NAME 

SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POINT 

Phosphorus  
Arsenic  

P 

As 

31.04 

74.96 

1.83 
5.73 

44° 

sublimes 

Antimony  
-Bismuth 

Sb 
Bi 

120.2 

208 

6.52 

9  80 

630° 

271° 

The  family.  The  elements  constituting  this  family  be- 
long in  the  group  with  nitrogen,  and  therefore  resemble  it  in 
a  general  way.  They  exhibit  a  regular  gradation  of  prop- 
erties, as  is  shown  in  the  above  table.  The  same  general  gra- 
dation is  also  found  in  their  chemical  character,  phosphorus 
being  an  acid-forming  element,  while  bismuth  is  essentially  a 
metal.  The  other  two  elements  are  intermediate  in  character. 


PHOSPHORUS 

History.  The  element  phosphorus  was  discovered  by  the 
alchemist  Brandt,  of  Hamburg,  in  1669,  while  searching  for 
the  philosopher's  stone  (p.  12).  Owing  to  its  peculiar 
properties  and  the  secrecy  which  was  maintained  about  its 
preparation,  it  remained  a  very  rare  and  costly  substance 
until  the  demand  for  it  in  the  manufacture  of  matches 
brought  about  its  production  on  a  large  scale. 

Occurrence.  Phosphorus  occurs  almost  entirely  in  vari- 
ous mineral  forms  of  calcium  phosphate.  Phosphorite  is  the 
most  abundant  of  these  minerals,  while  apatite  consists  of 

247 


248 


FIRST  COUESE  IN  CHEMISTEY 


calcium  phosphate,  together  with  calcium  fluoride  or  chloride. 
Calcium  phosphate  is  the  chief  mineral  constituent  of  the 
bones  of  animals,  and  bone  ash  is  therefore  nearly  pure 
calcium  phosphate. 

Preparation.  Phosphorus  is  manufactured  from  mineral 
phosphate  by  heating  the  phosphate  with  sand  and  carbon  in 
an  electric  furnace.  The  reaction  may  be  represented  in  two 
steps  : 


0  +  3Si02 
2P205+10C 


3  CaSi03+  P20t 
10CO  +  P, 


The  materials  are  fed  in  at  A  (Fig.  118)  by  the  feed  screw  B. 
The  phosphorus  vapor  escapes  at  D  and  is  condensed  under 

water,  while  the  calcium  sili- 
cate is  tapped  off  as  a  liquid 
at  C.  The  phosphorus  obtained 
in  this  way  is  quite  impure, 
and  is  purified  by  distillation. 

Properties.  The  purified 
phosphorus,  called  white  or 
yellow  phosphorus,  is  a  nearly 
colorless,  translucent,  waxy 
solid  which  melts  at  44°  and 
boils  at  287°.  It  can  there- 
fore be  cast  into  any  conven- 
ient form  under  warm  water, 
and  is  usually  sold  011  the 
market  in  the  form  of  sticks 


FIG. 


118.    A  furnace  for  the  pro- 
duction of  phosphorus 


(Fig.  119).  It  can  be  cut  with  a  knife,  but  this  must  always 
be  done  under  water,  since  phosphorus  is  extremely  inflam- 
mable, and  the  friction  of  the  knife  blade  is  almost  sure  to 
set  it  on  fire  if  cut  in  the  air.  It  is  not  soluble  in  water, 
but  is  freely  soluble  in  some  other  liquids,  notably  in  carbon 
disulfide.  Its  density  is  1.8. 


THE  PHOSPHORUS  FAMILY 


249 


Chemical  conduct.  When  exposed  to  the  air,  phosphorus 
slowly  combines  with  oxygen,  and  in  so  doing  gives  out  a 
pale  light,  or  phosphorescence,  which  can  be  seen  only  in  a 
dark  place.  The  heat  of  the  room  may  raise  the  tem- 
perature of  phosphorus  to  the  kindling  point,  when  it  burns 
with  a  sputtering  flame,  giving  off  dense  fumes  of  oxide  of 
phosphorus.  It  burns  with  dazzling  brilliancy  in  oxygen, 
and  combines  directly  with  many 
other  elements.  On  account  of  its 
great  attraction  for  oxygen,  it  is 
preserved  under  water. 

Phosphorus  is  very  poisonous, 
from  0.2  to  0.3  g.  being  a  fatal 
dose.  Ground  with  flour  and  grease 
or  similar  substances,  it  is  used  as 
a  poison  for  rats  and  other  vermin. 

Red  phosphorus.  On  standing, 
white  phosphorus  gradually  under- 
goes a  remarkable  change,  being 
converted  into  a  dark-red  powder 
which  has  a  density  varying  from  2.1 
to  2.38  and  is  called  red  phosphorus. 
It  no  longer  takes  fire  easily,  nor 
is  it  soluble  in  carbon  disulfide. 

It  is  not  poisonous  and,  in  fact,  is  an  entirely  different  sub- 
stance. The  velocity  of  this  change  of  white  phosphorus  to 
red  phosphorus  increases  with  rise  in  temperature,  and  red 
phosphorus  is  therefore  prepared  by  heating  the  white  form 
a  little  below  the  boiling  point.  When  distilled  and  quickly 
condensed,  the  red  form  changes  back  to  the  white. 

Matches.  The  chief  use  of  phosphorus  is  in  the  manu- 
facture of  matches,  two  general  varieties  of  which  are  in  com- 
mon use.  Ordinary  friction  matches  are  made  by  dipping 


FIG.  119.    Sticks  of  white 
phosphorus 


250  FIRST  COURSE  IN  CHEMISTRY 

the  match  sticks  first  into  some  inflammable  substance, 
such  as  melted  paraffin,  and  afterward  into  a  paste  consist- 
ing of  (1)  phosphorus  sesquisulfide  (P4S8),  (2)  some  oxidiz- 
ing substance,  such  as  manganese  dioxide  or  potassium 
chlorate,  and  (3)  a  binding  material,  usually  some  kind  of 
glue.  The  phosphorus  sulfide  is  ignited  by  friction,  the 
combustion  being  sustained  by  the  oxidizing  agent  and 
communicated  to  the  wood  by  the  burning  paraffin.  In 
sulfur  matches  the  paraffin  is  replaced  by  sulfur. 

In  safety  matches,  a  mixture  of  red  phosphorus,  an  oxi- 
dizing agent,  and  some  gritty  material,  such  as  powdered 
glass,  is  placed  on  the  side  of  the  box,  while  the  match  tip 
is  provided  with  an  oxidizing  agent  and  an  easily  oxidizable 
substance,  usually  antimony  sulfide.  The  match  cannot  be 
ignited  easily  by  friction  except  on  the  prepared  surface. 

Matches  were  formerly  made  from  white  phosphorus,  and 
the  workmen  frequently  suffered  from  dreadful  diseases  of  the 
bones  of  the  face.  On  this  account  the  manufacture  and  use 
of  matches  containing  white  phosphorus  was  gradually  pro- 
hibited in  European  countries.  It  was  found  that  the  com- 
pound P4S3,  which  is  easily  prepared  from  white  phosphorus, 
serves  just  as  well,  and  does  not  occasion  disease,  and  in  1913 
the  government  of  the  United  States  placed  a  prohibitive  tax 
(two  cents  per  hundred  matches)  on  the  white  phosphorus 
match,  at  the  same  time  forbidding  both  the  import  and 
export  of  such  matches. 

Phosphine  (PH3).  Phosphine  is  usually  made  by  heating 
phosphorus  with  a  solution  of  potassium  hydroxide,  the 
reaction  being  a  complicated  one. 

A  concentrated  solution  of  potassium  hydroxide,  together 
with  several  small  bits  of  phosphorus,  is  placed  in  the  flask  A 
(Fig.  120)  and  a  current  of  coal  gas  is  passed  into  the  flask 
through  the  tube  B  until  all  the  air  has  been  displaced.  The 


THE  PHOSPHORUS  FAMILY 


251 


gas  is  then  turned  off  and  the  flask  is  heated.  Phosphine  is 
formed  in  small  quantities  and  escapes  through  the  delivery 
tube,  the  exit  of  which  is  just  covered  by  the  water  in  the 
vessel  C:  Each  bubble  of  the  gas  as  it  escapes  into  the  air 
takes  fire,  and  the  product  of  combustion  (P20g)  forms  beau- 
tiful rings,  which  float  unbroken  for  a  time  in  quiet  air. 

Properties.  Phosphine  is  a  gas  of  unpleasant  odor  and  is 
exceedingly  poisonous.  Like  ammonia,  it  forms  salts  with 
the  halogen  acids.  Thus, 
we  \\avephosphoninm  chlo- 
ride (PH4C1)  analogous 
to  ammonium  chloride 
(NH.Cl). 

Oxides  of  phosphorus. 
Phosphorus  forms  two 
well-known  oxides  —  the 
trioxide  (P2O3)  and  the 
peiitoxide  (P2O5)  (some- 
times called  phosphoric 
anhydride).  The  pentox- 
ide  is  much  the  better 
known  of  the  two.  It  is  a 
snow-white,  voluminous 

powder  whose  most  marked  property  is  its  great  attraction 
for  water.  It  has  no  chemical  action  upon  most  gases,  so  that 
they  can  be  very  thoroughly  dried  by  being  passed  through 
properly  arranged  vessels  containing  phosphorus  pentoxide. 

Phosphoric  acid  (H3POJ.  This  acid  can  be  obtained  by 
dissolving  phosphorus  pentoxide  in  boiling  water,  but  it  is 
usually  made  by  treating  calcium  phosphate  with  concen- 
trated sulfuric  acid  : 


FIG.  120.  The  preparation  of  pliosphine 


Ca3(P04)2+  3 


3  CaSO4  +  2  H3PO4 


252  FIRST  COURSE  IN  CHEMISTRY 

The  calcium  sulfate  produced  in  the  reaction  is  nearly 
insoluble,  and  can  be  filtered  off,  leaving  the  phosphoric 
acid  in  solution.  This  acid  forms  large,  colorless  crystals 
which  are  exceedingly  soluble  in  water.  Being  a  tribasio 
acid,  it  forms  acid  as  well  as  normal  salts.  Thus,  the 
following  compounds  of  sodium  are  known : 

NaH2PO4    ....     primary,  or  mono  sodium-hydrogen  phosphate 
Na2HPO4    ....     secondary,  or  disodium-hydrogen  phosphate 
Na3PO4 tertiary,  or  normal  sodium  phosphate 

They  may  be  prepared  by  bringing  together  phosphoric  acid 
and  appropriate  quantities  of  sodium  hydroxide.  Phosphoric 
acid  also  forms  mixed  salts ;  that  is,  salts  containing  two  dif- 
ferent metals.  The  most  familiar  compound  of  this  kind  is 
microcosmic  salt,  which  has  the  formula  Na(NH4)HPO4. 

Phosphates.  The  phosphates  form  an  important  class  of 
salts.  The  normal  salts  are  nearly  all  insoluble,  and  many 
of  them  occur  in  nature.  The  secondary  phosphates  are  as 
a  rule  insoluble,  while  most  of  the  primary  salts  are  soluble. 
The  most  important  phosphate  is  calcium  phosphate,  which  is 
mined  in  enormous  quantities  for  use  as  a  fertilizer  (p.  309). 

OTHER  IMPORTANT  COMPOUNDS  OF  PHOSPHORUS 

Phosphorous  acid  (H3PO;;)  :  transparent,  colorless  crystals 
Metaphosphoric  acid  (PIPO3)  :  a  white  solid,  sold  in  sticks 
Pyrophosphoric  acid  (H4P2O7) :  a  white,  crystalline  solid 
Phosphorus  trichloride  (PC13)  :  a  colorless,  fuming  liquid 
Phosphorus  pentachloride  (PC15)  :  a  white,  fuming  solid 
Phosphorus  sesquisulfide  (P4S3)  :  a  solid  used  in  making  matches 

ARSENIC 

Occurrence.  Arsenic  occurs  in  nature  as  the  native  ele- 
ment, as  the  sulfides  realgar  (As2S2)  and  orpiment  (As0S8), 
as  the  oxide  As2Og,  and  as  a  constituent  of  many  metallic 
sulfides,  of  which  arsenopyrite  (FeAsS)  is  the  best  known. 


THE  PHOSPHOEUS  FAMILY  253 

Preparation.  The  element  is  prepared  by  heating  the 
arsenopyrite  in  iron  tubes,  out  of  contact  with  air,  when 
the  reaction  expressed  by  the  following  equation  occurs : 

4  FeAsS >-  4  FeS  +  As4 

The  arsenic,  being  volatile,  condenses  in  chambers  con- 
nected with  the  heated  tubes.  It  is  also  made  from  the 
oxide  by  reduction  with  carbon: 

2  As203  +  3  C >•  As4  +  3  CO2 

Properties.  Arsenic  is  a  steel-gray  solid  which  resembles 
coke  in  appearance  and  is  somewhat  brittle.  When  strongly 
heated,  it  sublimes  (that  is,  it  passes  into  a  vapor  without 
melting),  and  condenses  again  to  a  crystalline  solid  when 
the  vapor  is  cooled.  About  0.5  per  cent  of  arsenic  is 
added  to  lead  which  is  to  be  used  for  making  shot,  since 
it  greatly  increases  the  hardness  of  the  lead.  Unlike  most 
of  its  compounds,  the  element  itself  is  not  poisonous,  owing 
to  the  fact  that  it  is  not  soluble  in  the  digestive  fluids. 

Arsine  (AsH3).  When  any  compound  containing  arsenic 
is  brought  into  the  presence  of  nascent  hydrogen,  arsine 
(AsHg)  is  formed.  The  reaction,  when  oxide  of  arsenic  is 
so  treated,  is  represented  by  the  equation 

As2°3  +  12  [H]  =  2  AsH3  +  3  H20 

The  symbol  [H]  indicates  that  the  hydrogen  must  be  in 
the  nascent  state. 

Arsine  is  a  gas  with  a  peculiar,  garliclike  odor,  and  is 
intensely  poisonous.  It  is  an  unstable  compound,  decompos- 
ing into  its  elements  when  heated  to  a  moderate  temperature. 
It  is  combustible,  burning  with  a  pale  bluish- white  flame 
to  form  arsenic  trioxide  and  water  when  air  is  in  excess : 

2  AsH3  +  3  O2  — »•  As2O3  +  3  H2O 


254 


FIRST  COURSE  IN  CHEMISTRY 


When  the  supply  of  air  is  deficient,  water  and  metallic 
arsenic  are  the  products : 

4  AsH3  +  3  O2 >-  6  H2O  +  As4 

These  reactions  make  the  detection  of  even  minute  quan- 
tities of  arsenic  a  very  easy  matter. 

Marsh's  test  for  arsenic.  The  method  devised  by  Marsh  for 
detecting  arsenic  is  most  frequently  used,  the  apparatus  being 
shown  in  Fig.  121.  Hydrogen  is  generated  in  the  flask  A  by 
the  action  of  dilute  sulf  uric  acid  on  zinc,  is  dried  by  being  passed 

over  calcium  chlo- 
ride in  the  tube  B, 
and,  after  passing 
through  the  hard- 
glass  tube  C,  is  ig- 
nited at  the  jet  D. 
If  a  substance  con- 
taining arsenic  is 
now  introduced  into 
the  generator  A ,  the 
arsenic  is  converted 
into  arsine  by  the 
action  of  the  nascent 
hydrogen,  and  passes  to  the  jet  along  with  the  hydrogen.  If 
the  tube  C  is  strongly  heated  at  some  point  near  the  middle, 
the  arsine  is  decomposed  while  passing  this  point  and  the 
arsenic  is  deposited  just  beyond  the  heated  point  in  the  form 
of  a  shining  brownish-black  mirror.  A  small  fraction  of  a 
milligram  of  arsenic  can  be  detected  by  this  test. 

Arsenious  oxide  (white  arsenic)  (As203).  Arsenious  oxide 
is  obtained  as  a  by-product  in  various  metallurgical  proc- 
esses, and  in  burning  pyrite  which  contains  arsenopyrite 
for  the  production  of  sulfur  dioxide : 

2  FeAsS  4-  5  00 ^  Fe  Oa  +  As  Oa  +  2  SO2 


FIG.  121.    Marsh's  apparatus  for  the  detection 
of  arsenic 


THE  PHOSPHORUS  FAMILY  255 

The  arseniotis  oxide  is  condensed  in  appropriate  chambers. 
It  is  a  rather  heavy  substance,  obtained  either  as  a  crystal- 
line powder  or  in  lumps  resembling  porcelain  in  appearance. 
It  is  very  poisonous,  from  0.2  to  0.3  g.  being  a  fatal  dose. 
Arsenious  oxide  is  used  in  glass-making  and  in  the  dye 
industry,  and  as  the  source  from  which  all  other  arsenic 
compounds  are  made. 

Arsenic  insecticides.  Several  compounds  of  arsenic  have 
important  uses  as  insecticides.  Paris  green  and  8cheeles 
green  are  made  by  treating  solutions  of  copper  salts  with 
arsenious  oxide.  Lead  arsenate  (Pbg(AsO4)2)  is  widely 
used  in  connection  with  lime-sulfur  sprays  (p.  144). 

OTHER  COMPOUNDS"  OF  ARSENIC 

Arsenious  acid  (H3AsO3)  :  known  only  in  solution 

Arsenic  acid  (H3AsO4)  :  colorless  transparent  crystals 

Arsenious  sulfide  (As2S3)  :  a  yellow  solid,  prepared  by  treating  an 

arsenious  compound  with  hydrogen  sulfide 
Arsenic  sulfide  (As2S5)  :  a  yellow  solid,  prepared  by  treating  an  arsenic 

compound  with  hydrogen  sulfide 

ANTIMONY 

Occurrence.  Antimony  occurs  in  nature  chiefly  as  the 
sulfide  Sb2S3,  called  stibnite,  though  it  is  also  found  as 
oxide  and  as  a  constituent  of  many  complex  minerals. 

Preparation.  Antimony  is  prepared  from  the  sulfide  in 
a  very  simple  manner.  The  sulfide  is  melted  with  iron 
in  a  furnace,  when  the  iron  combines  with  the  sulfur  to 
form  a  liquid  layer  of  melted  iron  sulfide,  while  the  heavier 
antimony  sinks  to  the  bottom  and  is  drawn  oft7  from  time  to 
time.  The  reaction  involved  is  represented  by  the  equation 

Sb  S,  4-  3  Fe  — >-  2  Sb  +  3  FeS 

2     3 


256  FIRST  COUKSE  IN  CHEMISTEY 

Properties.  Antimony  is  a  silvery,  metallic  substance 
whose  density  is  6.52.  It  is  highly  crystalline,  hard,  and 
very  brittle.  For  a  metal  it  has  a  rather  low  melting  point 
(630°),  and  it  expands  very  noticeably  on  solidifying.  Its 
chief  use  is  in  making  Babbitt  metal  and  other  alloys 
(p.  258). 

Chemical  conduct.  In  chemical  conduct,  antimony  re- 
sembles arsenic  in  many  particulars.  It  forms  the  oxides 
Sb2O3  and  Sb2O5  and  the  hydride  SbH3  (called  stiUne).  It 
also  forms  a  number  of  acids.  It  is,  however,  much  more 
pronouncedly  a  metal,  dissolving  in  concentrated  sulfuric 
acid  to  form  the  sulfate  Sb2(SO4)3,  and  in  aqua  regia  to  form 
the  chloride  SbCl3.  Hydrogen  sulfide  passed  into  a  solution 
of  the  chloride  precipitates  the  orange-colored  sulfide  Sb2S3, 
having  the  same  composition  as  the  black  mineral  stibnite. 

BISMUTH 

Occurrence  and  preparation.  Bismuth  is  usually  found 
in  nature  in  the  uncombined  form.  It  also  occurs  as  oxide 
and  sulfide.  Most  of  the  bismuth  of  commerce  comes  from 
Saxony  and  from  Australia  and  Colorado,  but  it  is  not  an 
abundant  element.  The  metal  is  prepared  by  merely  heating 
an  ore  containing  the  native  bismuth,  and  allowing  the 
melted  metal  to  run  out  into  suitable  vessels.  Other  ores  are 
converted  into  oxides  and  reduced  by  heating  with  carbon. 

Properties.  Bismuth  is  a  heavy,  crystalline,  brittle  metal 
nearly  the  color  of  silver,  but  with  a  slightly  rosy  tint  which 
distinguishes  it  from  other  metals.  It  melts  at  a  low  tem- 
perature (271°)  and  has  a  density  of  9.8.  It  is  not  acted 
upon  by  hydrochloric  acid,  being  below  hydrogen  in  the 
displacement  series  (p.  112),  but  nitric  and  sulfuric  acids 
act  upon  it  in  the  same  way  that  they  do  upon  copper. 


THE  PHOSPHORUS  FAMILY  257 

Compounds  of  bismuth.  Unlike  the  other  elements  of 
this  group,  bismuth  has  almost  no  acid  properties.  Its  chief 
oxide,  Bi2Og,  is  basic  in  its  properties.  It  dissolves  in  con- 
centrated acids  and  forms  salts  of  bismuth : 

Bi203  +  6  HC1  =  2  BiCl3  +  3  H2O 
Bi203  +  6  HN03  =  2  Bi(N03)3  +~3  H2O 

The  nitrate  and  chloride  of  bismuth  can  be  obtained  as 
well-formed  colorless  crystals.  When  treated  with  water 
the  salts  are  decomposed  in  the  manner  explained  in  the 
following  paragraph : 

Hydrolysis  of  bismuth  salts.  The  hydroxide  Bi(OH)3 
is  a  very  weak  base,  and  we  should  expect  its  salts  to  be 
decomposed,  or  hydrolyzed,  by  water  (p.  138).  If  bismuth 
chloride  (BiCl8),  were  to  be  completely  hydrolyzed,  the 
equation  would  be  as  follows : 

/Cl  /OH 

Bif-Cl  +  3  II2O ^Bi^OII  +  3  HC1 

\C1  \OH 

The  reaction  is  not  so  complete,  however  —  only  two  of 
the  three  chlorine  ions  being  replaced  by  hydroxyl  ions : 

/Cl  /Cl 

Bi^-Cl  +  2  II0O hBi^-OH  +  2  IIC1 

\C1  \OH 

If  we  wish  to  prevent  this  hydrolysis,  we  must  add  hydro- 
chloric acid  in  sufficient  quantity  to  reverse  the  reaction  of 
hydrolysis  (p.  136).  The  compound  Bi(OH)2Cl  easily  loses 
water,  forming  a  white  solid,  BiOCl,  known  as  Usmuthyl 
chloride.  A  similar  compound,  BiONO3,  is  known  as  bis- 
muth subnitrate,  and  is  used  largely  in  medicine. 

Basic  salts.  The  compound  formed  by  the  partial  hydrol- 
ysis of  bismuth  chloride  is  unlike  any  AVC  have  yet  met. 
Since  it  contains  hydroxyl  radicals  combined  with  a  metal, 


258 


FIRST  COURSE  IN  CHEMISTRY 


we  must  regard  it  as  a  base ;  but  it  also  contains  a  chlorine 
atom  combined  with  a  metal,  so  that  it  is  likewise  a  salt. 
Since  it  has  the  characteristics  of  both  a  base  and  a  salt,  it 
is  called  a  basic  salt.  A  basic  salt  may  in  general  be  regarded 
as  a  base,  a  part  of  whose  hydroxyl  groups  have  been  replaced 
by  an  acid: 


/OH 
Bi-OH  +  IIC1 


Bi^-OH  +  H20 

\OH 


base 


acid 


basic  salt 


water 


ALLOYS 

Some  metals  when  melted  together  thoroughly  intermix, 
and  on  cooling  form  a  metallic-appearing  substance  called  an 
alloy.  Not  all  metals  will  mix  in  this  way,  and  in  some  cases 


FIG.  122.   An  automatic  fire  curtain  over  a  door 

definite  chemical  compounds  are  formed  and  separate  out  as 
the  mixture  solidifies,  thus  destroying  the  uniform  quality 
of  the  alloy.  In  general,  the  melting  point  of  the  alloy  is 
below  the  average  of  the  melting  points  of  its  constituents, 
and  it  is  usually  lower  than  that  of  any  one  of  them. 

Both  antimony  and  bismuth  alloy  readily  with  many  other 
metals.  The  alloys  so  formed  are  heavy,  are  easily  melted, 
do  not  oxidize  easily  nor  act  upon  water,  and,  in  general,  are 


THE  PHOSPHORUS  FAMILY  259 

well  adapted  to  many  technical  uses.  The  manufacture  of 
alloys  constitutes  the  chief  use  of  these  two  metals. 

Antimony  imparts  to  its  alloys  the  property  of  expanding 
slightly  in  solidification,  which  renders  them  especially  use- 
ful in  type  founding,  where  fine  lines  are  to  be  reproduced 
on  a  cast.  Type  metal  consists  of  antimony,  lead,  and  tin. 
Babbitt  metal,  used  for  journal  bearings  in  machinery,  con- 
tains the  same  metals  in  a  different  proportion,  together 
with  a  small  percentage  of  copper. 

Bismuth  is  particularly  valuable  in  the  production  of 
very  low-melting  alloys.  For  example,  Wood's  metal,  con- 
sisting of  bismuth,  lead,  tin,  and  cadmium,  melts  at  60.5°. 
The  low  melting  point  of  such  alloys  is  turned  to  practical 
account  in  making  safety  plugs  in  boilers,  automatic  fire 
curtains  and  automatic  water  sprinklers  in  buildings,  and 
many  similar  devices. 

Fig.  122  shows  a  fire  curtain,  which  is  held  in  place  by  two 
wires  A,A}  joined  at  B  by  a  bismuth  alloy.  In  case  of  fire,  the 
alloy  melts  and  the  wires  holding  the  curtain  up  are  thereby 
released  and  the  curtain  drops,  covering  the  door. 

EXERCISES 

1.  Xame  all  the  elements  so  far  studied  which  may  be  obtained 
in  different  forms. 

2.  In  the  preparation  of  phosphine,  why  is  coal  gas  passed  into 
the  flask  ?    What  other  gases  would  serve  the  same  purpose  ? 

3.  Could  phosphoric  acid  be  substituted  for  sulfuric  acid  in  the 
preparation  of  the  common  acids  ? 

4.  What  weight  of  arsenic  can  be  obtained  by  heating  500  g.  of 
arsenopyrite ?    Ans.  230.12  g. 

5.  What  was  the  philosopher's  stone? 

TOPICS  FOR  THEMES 

The  invention  of  matches  (see  encyclopedia). 
Alloys  (see  encyclopedia). 


CHAPTER  XXVII 
SILICON  AND  BORON 

SILICON 

Occurrence.  Next  to  oxygen,  silicon  is  the  most  abundant 
element.  It  occurs  in  nature  chiefly  in  the  form  of  SiO  , 
called  silicon  dioxide  (silica) ;  or  in  the  form  of  salts  of  silicic 
acids  (silicates).  These  compounds  form  a  large  fraction  of 


FIG.  123.  Carborundum  crystals  and  abrasive  objects  made  of  carborundum 

the  earth's  crust,  constituting  almost  all  of  the  common 
rocks  except  limestone.  Plants  absorb  small  amounts  of 
silica  from  the  soil,  and  it  is  also  found  in  minute  quantities 
in  animal  organisms,  especially  in  hair,  claws,  and  horns. 

260 


SILICON  AND  BORON  261 

The  element.  Silicon  is  prepared  on  the  large  scale  by 
heating  pure  sand  with  carbon  in  an  electric  furnace  to  a 
very  high  temperature  : 

Si0  +  2C  —  ^Si  + 


It  is  a  silvery,  metallic-appearing  substance,  highly  crystal- 
line and  very  brittle.  It  is  used  in  the  manufacture  of 
certain  varieties  of  iron. 

Silicides.  As  the  name  indicates,  silicides  are  compounds 
consisting  of  silicon  and  some  one  other  element.  They  are 
very  stable  at  high  temperatures,  and  are  usually  made  by 
heating  the  appropriate  substances  in  an  electric  furnace. 

The  most  important  silicide  is  car- 
borundum, which  is  a  silicide  of  car- 
bon  of  the  formula  CSi.  It  is  made 
by  heating  coke  and  sand  in  an 
electric  furnace,  the  process  being 
extensively  carried  on  at  Niagara 
Falls.  The  following  equation  rep-  FlG  124  Cross  section 

resents   the   reaction  :  of     a    charged     carbo- 

rundum furnace  before 
SiO0  +  3  C  -  >  CSi  +  2  CO  heating 


2 


The  substance  so  prepared  consists  of  beautiful  purplish- 
black  crystals,  which  are  surpassed  in  hardness  only  by  the 
diamond  and  by  boron  carbide.  Carborundum  is  used  as  an 
abrasive ;  that  is,  as  a  material  for  grinding  and  polishing 
very  hard  substances.  Fig.  123  shows  two  samples  of  the 
crystalline  material,  as  well  as  whetstones  and  a  grinding- 
wheel  prepared  from  carborundum. 

Manufacture  of  carborundum.  The  mixture  of  materials  is 
heated  in  a  large  resistance  furnace,  similar  to  the  one  em- 
ployed in  the  manufacture  of  graphite  (p.  195).  Fig.  124  rep- 
resents a  cross  section  of  the  furnace  after  charging,  A  being 


262 


FIRST-  COURSE  IN  CHEMISTRY 


FIG.  125.    Cross  section 
of    a    charged    carbo- 
rundum furnace  after 
heating 


the  carbon  core  and  B  the  coke  and  sand.  Fig.  125  shows  the 
appearance  after  heating.  A  is  the  core  of  carbon,  surrounded 
by  crystallized  carborundum  B.  Around 
this  is  a  shell  of  amorphous  carborundum 
C,  while  D  is  unchanged  charge. 

Silicon  dioxide  (silica)  (Si02).  This 
substance  is  found  in  a  great  variety 
of  forms  in  nature,  both  in  the  amor- 
phous and  in  the  crystalline  condition. 
In  the  form  of  quartz  (Fig.  126)  it  is 
found  in  beautifully  formed  six-sided 
prisms,  sometimes  of  great  size.  When 
pure  it  is  perfectly  transparent  and  colorless.  Some  colored 
varieties  are  given  special  names,  as  amethyst  (violet), 
rose  quartz  (pale  pink),  smoky  or  milky  quartz  (colored 
and  opaque).  Other 
varieties  of  silicon  di- 
oxide, some  of  which 
also  contain  water, 
are  chalcedony,  onyx, 
jasper,  opal,  agate, 
and  flint.  Sand  and 
sandstone  are  largely 
silicon  dioxide. 

The  skeletons  of 
certain  microorgan- 
isms (infusoria)  are 
composed  of  nearly 

pure  silica.  In  some  ^  m  Quartz  crysta]s 

localities  these  have 

accumulated  in  immense  deposits,  forming  a  very  fine  and 
sharp  sand,  called  infusorial  earth.  This  material  is  often 
used  as  a  scouring  powder,  especially  in  scouring-soaps. 


SILICON  AND  BORON 


263 


Properties.  As  obtained  by  chemical  processes,  silicon 
dioxide  is  an  amorphous  white  powder.  In  the  crystallized 
state  it  is  very  hard,  and  has  a  density  of  2.6.  Pure  silica 
begins  to  soften  at  about  1600°,  and  somewhat  above  this 
temperature  it  can  be  drawn  out  into  threads  and  blown 
into  tubes  and  small  vessels  like  glass.  These  articles  are 
attacked  by  comparatively 
few  ordinary  reagents,  and 
they  do  not  expand  or  con- 
tract to  any  appreciable  ex- 
tent with  even  very  great 
changes  in  temperature.  On 
this  account  a  quartz  vessel 
can  be  heated  red  hot  and 
plunged  into  cold  water  with- 
out cracking.  Fig.  127  shows 
a  quartz  crucible  and  quartz 
tubes  on  a  wire  triangle  used 
to  support  the  crucible  when 
heated. 

Chemical  conduct.  Silica 
is  insoluble  in  water  and  in 
most  acids.  It  is  very  stable, 
so  that  the  oxygen  which  it  contains  can  be  removed  only 
by  the  most  powerful  reducing  agents,  and  at  very  high 
temperatures.  Hydrofluoric  acid  attacks  it  readily  (p.  168), 
according  to  the  equation 

Si02  +  2  H2F2  — *  SiF4  +  2  HaO 

Since  it  is  the  anhydride  of  an  acid,  it  dissolves  in  fused 
alkalies  to  form  silicates.  Being  nonvolatile,  it  will  drive 
out  most  other  anhydrides  when  heated  to  a  high  tem- 
perature with  their  salts,  especially  when  the  silicates  so 


FIG.  127.  A  crucible  and  a  triangle 
made  from  quartz 


264  FIKST  COURSE  IN  CHEMISTRY 

formed  are  fusible.    The  following  equations  illustrate  this 


.  — ^Na2SiO3  +  CO0 
Na'SO."  +  Si02  — >•  NaoSiOs  +  SO, 

Sv  A  Jo  o 

Silicic  acids.  Silicon  forms  two  simple  acids,  orthosilicic 
acid  (H4SiO4)  and  metasilicic  acid  (H2SiO3).  Orthosilicic 
acid  is  formed  as  a  jellylike  mass  (colloid)  when  ortho- 
silicates  are  treated  with  strong  acids.  If  one  attempts  to 
dry  this  acid,  it  loses  water,  passing  into  metasilicic  acid 
(common  silicic  acid)  : 

H4Si04^H2Si03  +  H20 

Metasilicic  acid,  when  heated,  breaks  up  into  silica  and 
water,  thus: 

H2Si03 >•  H20  +  Si02 

Salts  of  silicic  acids  ;  silicates.  A  number  of  salts  of  the 
orthosilicic  and  metasilicic  acids  occur  in  nature.  Thus, 
mica  (KAlSiO4)  is  a  mixed  salt  of  orthosilicic  acid. 

Polysilicic  acids.  Silicon  has  the  power  to  form  a  great 
many  complex  acids  which  may  be  regarded  as  derived 
from  the  union  of  several  molecules  of  orthosilicic  acid,  with 
the  loss  of  water.  These  are  called  polysilicic  acids.  For 
example,  we  have 

3  H4Si04  — >-  H4Si308  +  4  H20 

Salts  of  these  acids  make  up  the  great  bulk  of  the  earth's 
crust.  Feldspar,  for  example,  has  the  formula  KAlSi3Og, 
and  is  a  mixed  salt  of  the  acid  H4Si3Og,  whose  formation 
is  represented  in  the  equation  above.  ^Kaolin,  or  pure  clay, 
has  the  formula  H4Al0Si209,  or  as  commonly  written, 
Al2Si2O7  •  2  H2O.  Granite  is  composed  of  crystals  of  feld- 
spar and  mica,  cemented  together  with  amorphous  silica. 


SILICON  AND  BOEON  265 

Water  glass.  A  concentrated  solution  of  sodium  silicate 
(Na2SiO3)  or  of  potassium  silicate  (K2SiO3)  or  of  both  is 
called  water  glass.  It  is  a  thick,  sticky  liquid  made  by 
fusing  sand  with  the  carbonate  of  sodium  or  of  potassium. 
It  is  used  for  the  purpose  of  giving  a  glazed  waterproof 
surface  to  porous  materials,  such  as  wood,  stone,  and  plas- 
ter ;  to  render  curtains  noninflammable  ;  as  a  glue  for  glass 
and  pottery ;  and  as  an  ingredient  in  cheap  soaps. 

Its  property  of  closing  pores  is  turned  to  account  in  preserv- 
ing eggs  for  winter  use.  The  eggs  are  packed  in  crocks  and 
then  covered  with  a  liquid  made  by  adding  1  volume  of  com- 
mercial water  glass  to  10  volumes  of  water.*  Over  the  liquid  is 
then  poured  a  little  melted  paraffin,  which  soon  hardens  and 
excludes  the  air.  Fresh  eggs  can  easily  be  preserved  for  from 
eight  to  ten  months  in  this  way. 

Glass.  When  sodium  silicate,  calcium  silicate,  and  silicon 
dioxide  are  heated  together  to  a  very  high  temperature, 
the  mixture  slowly  fuses  to  a  transparent  liquid,  which,  on 
cooling,  passes  into  the  rigid  material  called  glass.  Instead 
of  starting  with  the  silicates  of  sodium  and  calcium,  it  is 
more  convenient  and  economical  to  heat  sodium  carbonate 
(or  sulfate)  and  lime  with  an  excess  of  clean  sand,  the  sili- 
cates being  formed  during  the  heating : 

Na2C03  +  Si02 >•  Na2Si03  +  CO2 

CaO  +  SiO9 *•  CaSiO, 

Z  o 

Molding  and  blowing  glass.  The  way  in  which  the  melted 
mixture  is  handled  in  the  glass  factory  depends  upon  the  char- 
acter of  the  object  to  be  made.  Many  articles,  such  as  bottles, 
are  made  by  blowing  the  plastic  glass  into  hollow  molds  of  the 
desired  shape.  The  mold  is  opened,  a  lump  of  plastic  glass  on 
a  hollow  rod  is  lowered  into  it,  and  the  mold  is  then  closed. 
By  blowing  into  the  tube  the  glass  is  forced  into  the  shape  of 


266 


FIBST  COUKSE  IX  CHEMISTRY 


the  mold.  The  mold  is  then  opened  (Fig.  128)  and  the  object 
lifted  out.  The  top  of  the  object  must  be  cut  off  at  the  proper 
place  and  the  sharp  edges  rounded  off  in  a  flame.  Bottles 

are  now  more  often 
made  by  machinery,  in 
which  process  the  neck 
is  finished  first  and 
the  bottle  then  blown 
by  compressed  air. 

Other  objects,  such 
as  lamp  chimneys, 
glasses,  and  beakers, 
are  revolved  while  be- 
ing blown  in  the  mold, 
and  have  no  ridge 
showing  where  the 
mold  closes.  Window 


EIG.  128.   Making  a  glass  object  in  a  mold 


glass  is  made  by  gathering  a  lump  of  molten  glass  on  the  end 
of  a  hollow  rod  (Fig.  129)  and  blowing  this  into  the  form  of 
large  hollow  cylinders  (Fig.  130)  about  6  ft.  long  and  1-i-  ft. 
in  diameter.  These 
are  cut  longitudinally 
(Fig.  131),  and  are 
then  placed  in  an 
oven  and  heated  until 
they  soften,  when  they 
are  flattened  out  into 
plates  and  cut  into 
the  desired  sizes.  Sim- 
ilar cylinders  are  now 
made  by  dipping  a 
hollow  tube  into  the 
melted  glass  and  slowly 
withdrawing  it  while 
compressed  air  is  blown 
through  the  tube.  In  this  way  a  very  long  cylinder  is  formed. 
Plate  glass  is  cast  into  flat  slabs  (Fig.  132),  which  are  then 
ground  and  polished  (Fig.  133)  to  perfectly  plane  surfaces. 


FIG. 129.  F 


t  step  in  making  window  glass : 
blowing  the  ball 


SILICOX  AND  BORON 


267 


FIG.  130.    Second  step  in  making  window 
glass  :  blowing  the  cylinders 


Varieties  of  glass.  The  glass  made  from  sodium  carbonate, 
lime,  and  sand  is  soft  and  easily  fusible.  If  potassium  car- 
bonate is  substituted  for  the  sodium  carbonate,  the  glass  is 

much  harder  and  less 
easily  fused ;  increas- 
ing the  amount  of 
sand  has  somewhat  the 
same  effect.  Potassium 
glass,  of  which  Jena 
glass  is  a  variety,  is 
largely  used  in  making 
chemical  glassware, 
since  it  resists  the  ac- 
tion of  reagents  better 
than  the  softer  so- 
dium glass.  If  lead 
oxide  is  substituted 
for  the  whole  or  a 
part  of  the  lime,  the 

glass  is  very  soft,  but  has  a  high  index  of  refraction  and  is  valua- 
ble for  making  optical  instruments  and  artificial  jewels  (paste). 

Coloring  glass.  Va- 
rious substances  fused 
along  with  the  glass 
give  characteristic  col- 
ors. The  amber  color 
of  common  bottles  is 
due  to  iron  compounds 
in  the  glass  ;  in  other 
cases  iron  colors  the 
glass  green.  Cobalt 
compounds  color  it 
deep  blue ;  compounds 

of  manganese  ^  give  it        ^  lgL   Thh.d  step  in  making  window 
an  amethyst  tint,  and  glass .  cutting  the  cylinders 

uranium     compounds 

impart   a  peculiar   yellowish-green  color.    Iron   is   nearly  al- 
ways present  in  sand,  and  this  makes  a  green  glass  unless 


268 


FIRST  COURSE  IN  CHEMISTRY 


an  oxidizing  agent  is  used.  The  green  color  is  largely  removed 
by  the  addition  of  manganese  dioxide,  which  oxidizes  the  iron 
to  a  form  having  a  yellowish  tinge,  and  this  color  is  then  neu- 
tralized by  the  manganese,  since  the  yellow  produced  by  the 

iron  and  the  amethyst 
produced  by  the  man- 
ganese are  complemen- 
tary colors,  producing 
white  light. 
Nature  of  glass.  Glass 

jjK^'          »  *  sii  *s  no^  a  ^e^n^e  chem- 

B  ical  compound,  and 
its  composition  varies 
between  wide  limits. 
Fused  glass  is  really 


FIG.  132.    Casting  plate  glass 


a  solution  of  various 
silicates,  such  as  those 
of  calcium  and  lead,  in  fused  sodium  silicate  or  potassium 
silicate.  A  certain  amount  of  silicon  dioxide  is  also  present. 
This  solution  is  cooled  under  such  conditions  that  the  dissolved 
substances  do  not  separate  from  the  solvent.  The  compounds 
which  are  used  to  color  the  glass  are  sometimes  converted 
into  silicates  which 
then  dissolve  in  the 
glass,  giving  it  a  uni- 
form color.  In  other 
cases,  as  in  the  mi/ Jet/ 
glasses  which  resem- 
ble porcelain  in  ap- 
pearance, the  color  or 
opaqueness  is  due  to 
the  finely  divided  ma- 
terial evenly  distrib- 
uted throughout  the 

glass,  but  not  dissolved  in  it.  Milky  glass  is  made  by  mixing 
calcium  fluoride,  tin  oxide,  or  some  other  insoluble  substance 
in  the  melted  glass.  Selenium  or  gold  in  very  finely  divided 
(colloidal)  form  scattered  through  glass  gives  it  shades  of  red. 


FIG.  133.   Polishing  plate  glass 


SILICON  AND  BORON  269 

Enamels.  The  surface  of  metal  vessels,  such  as  cooking  uten- 
sils and  bath  tubs,  is  often  covered  by  a  kind  of  opaque  glass 
called  enamel  (granite  or  agate  ware).  This  contains  boric  ox- 
ide (B20g)  in  place  of  some  silica,  and  oxides  of  a  number  of 
different  metals,  such  as  barium,  zinc,  or  lead,  in  place  of  some 
of  the  calcium. 

BORON 

Occurrence.  Boron  occurs  in  nature  as  boric  acid  (HgBO3), 
and  in  salts  of  polyboric  acids,  which  usually  have  very  coin- 
plicated  formulas. 

Preparation  and  properties.  The  element  boron  is  ex- 
tremely difficult  to  prepare  in  pure  condition,  and  it  is  only 
known  in  an  amorphous  state.  Its  electrical  resistance 
varies  to  an  extraordinary  extent  with  changes  in  tempera- 
ture, and  this  property  promises  to  make  it  very  useful. 

Boric  acid  (H3B03).  This  compound  is  found  dissolved 
in  the  water  of  hot  springs  in  some  localities,  particularly 
in  Italy.  Being  volatile  with  steam,  boric  acid  is  present  in 
the  vapor  from  these  springs.  The  acid  is  easily  obtained 
from  these  sources  by  condensation  and  evaporation,  the 
necessary  heat  being  supplied  by  other  hot  springs. 

It  is  often  prepared  by  treating  a  strong,  hot  solution  of 
borax  (Na2B4O7)  with  sulfuric  acid.  Boric  acid,  being  but 
sparingly  soluble  in  water,  crystallizes  out  on  cooling : 

NaQB407  -f  5  H00  +  H2SO4  — >•  Na2SO4  +  4  H3BO3 

Z       ^        7  £  21  J  4  oo 

Boric  acid  crystallizes  in  pearly  flakes  which  are  slippery 
to  the  touch.  It  is  a  mild  antiseptic,  and  is  often  used  in 
medicine,  particularly  for  eyewashes.  Its  acid  properties 
are  extremely  weak. 

Metaboric  and  polyboric  acids.  When  boric  acid  is  gently 
heated,  it  is  converted  into  metaboric  acid  (HBO2)  : 

H3B03 


270 


FIRST  COURSE  IN  CHEMISTRY 


FIG.  134.    A  specimen  of 
colemanite 


On  heating  metaboric  acid  to  a  somewhat  higher  tempera- 
ture, tetraboric  acid  (H2B4O7)  is  formed : 

4  HBO2  — *  H2B4O7  +  H2O 

Borax.  The  sodium  salt  of  tetra- 
boric acid  has  the  formula  Na2B4O7. 
Borax  is  a  hydrate  (p.  271)  of  this 
salt.  It  is  found  in  some  arid  coun- 
tries, as  southern  California  and 
Tibet,  but  is 
now  made  com- 
mercially from 
the  mineral  cole- 
manite.  This 
is  the  calcium 
salt  of  a  com- 
plex boric  acid 

(Ca2B6On)  and  occurs  in  large  depos- 
its in  California.    When  it  is  treated 

with  a  solution  of  sodium  carbonate, 

calcium  carbonate  is  precipitated  and 

borax    crystallizes  from  the  solution. 

Fig.  134  shows  a  sample  of  coleman- 

ite,  and  Fig.  135  the  crystalline  borax 

obtained  from  it. 

When  heated,  borax  at  first  swells 

up  greatly,  and  then  melts  to  a  clear 

glass.   This  glass  has  the  property  of 

easily  dissolving  many  metallic  oxides, 

and  on  this  account  borax  is  used  as  a 

flux   in   brazing,  for  the    purpose  of   removing  from  the 

metallic  surfaces  to  be  brazed  the  film  of  oxide  with  which 

they  are  likely  to  be  covered  (p.  272).    These  oxides  often 


FIG.  135.    Crystals  of 
borax 


SILICON  AND  BORON  271 

give  a  characteristic  color  to  the  clear  borax  glass,  and  on 
this  account  borax  beads  are  often  used  in  testing  -for  the 
presence  of  metals. 

Borax  is  extensively  used  as  a  constituent  of  enamels 
and  glazes  for  both  metal  ware  and  pottery.  It  is  used 
to  soften  hard  water  for  domestic  purposes  (p.  300),  as 
a  mild  alkali  (like  soap),  as  an  antiseptic,  and  as  a  flux 
in  brazing. 

Hydrates ;  water  of  hydration.  The  water  with  which 
any  compound  unites  to  form  a  hydrate  is  termed  water  of 
hydration.  In  writing  the  formula  of  a  hydrate  it  is  cus- 
tomary to  represent  the  water  of  hydration  separately. 
Thus,  the  formula  of  borax  is  written  Na2B4O7  •  10  H2O. 
This  expresses  the  fact  that  1  molecule  of  borax  is  made 
up  of  1  molecule  of  sodium  tetraborate  (Na2B4O?)  and 
10  molecules  of  water  of  hydration.  The  hydrates  are  true 
chemical  compounds,  and  most  of  them  are  crystalline. 
When  a  hydrate  is  heated,  the  water  of  hydration  is  driven 
off.  A  salt  crystallizing  without  water  of  hydration  is  said 
to  be  anhydrous. 

EXERCISES 

1.  Account  for  the  fact  that  both  silicon  and  carborundum  can 
be  made  by  heating  sand  with  carbon. 

2.  Account  for  the  fact  that  a  solution  of  borax   in  water  is 
alkaline. 

3.  What  weight  of  water  of  hydration  does  1  kg.  of  borax  con- 
tain?   Am.  471  g. 

4.  What  weight  of  borax  can  be  made  from  a  ton  of  colemanite? 
Ans.    3559  Ib. 

5.  Why  does  not  sodium  silicate  form  a  glass  suitable  for  com- 
mon use  ? 

TOPICS  FOR  THEMES 

Glass-making  (Rogers  and  Aubert,  Industrial  Chemistry). 
Death  Valley  and  the  borax  beds  (see  encyclopedia). 


CHAPTER  XXVIII 
THE  METALS 

The  metals.  The  elements  so  far  considered  have  nearly 
all  been  those  whose  compounds  with  oxygen  and  hydrogen 
are  acids,  and  they  are  called  the  acid-forming  elements. 
Those  which  we  shall  now  study  are  known  collectively  as 
the  metals.  Their  hydroxides  are  bases,  and. on  this  account 
the  metals  are  sometimes  defined  as  those  elements  ivhose  hy- 
droxides are  bases.  When  the  hydroxide  of  a  metal  or  any  of 
the  simple  salts  derived  from  the  hydroxide  are  dissolved 
in  water,  the  metallic  element  forms  the  cation  and  carries 
a  positive  charge. 

Properties  of  the  metals.  The  metals  are  all  solids,  except 
mercury,  which  is  a  liquid.  Most  metals  have  a  high  den- 
sity, are  good  conductors  of  heat  and  electricity,  are  notably 
crystalline  in  structure,  and  take  a  bright  polish.  With  the 
exception  of  gold  and  copper,  they  have  a  silvery  luster. 
Most  of  them  combine  readily  with  oxygen  and  sulfur,  and 
their  surfaces  quickly  tarnish  on  exposure  to  the  air. 

A  few  of  the  least  active  of  the  metals,  such  as  gold  and 
copper,  occur  to  some  extent  in  nature  in  the  native  state. 
Most  of  them  are  found  in  combination  as  oxides,  as  sulfides, 
or  as  salts  of  various  acids,  especially  as  silicates.  The  proc- 
ess of  winning  metals  from  their  ores  is  called  metallurgy. 
The  details  of  the  metallurgy  of  any  given  metal  must  be 
adapted  to  the  properties  of  the  metal,  to  the  form  in  which 
it  is  found,  and  to  the  cost  of  the  operation. 

272 


THE  METALS  273 

Compounds  of  the  metals.  Since  the  metallic  elements  are 
base-forming  elements,  the  compounds  which  they  form  are 
chiefly  oxides,  hydroxides,  and  salts  of  various  acids.  We 
may  expect  each  metal  to  form  a  salt  with  each  acid,  and 
since  a  great  many  acids  are  known,  the  total  number  of 
salts  is  very  great.  Only  those  will  be  described  which 
have  important  uses  in  the  industries. 

Preparation  of  salts.  There  are  two  general  ways  of  pre- 
paring salts,  which  are  employed  so  often  that  it  is  well 
to  fix  them  in  mind  at  the  outset  of  the  study  of  metals. 

1.  Soluble  salts.    If  a  given  salt  is  soluble,  it  can  usually 
be  prepared  in  solution  by  treating  the  proper  metal,  or  its 
oxide,  hydroxide,  or  carbonate,  with  the  proper  acid.   All  of 
these  reactions  have  already  been  illustrated  repeatedly : 

Cu  +  2  H0S04  — >-  CuS04  +  S02  +  2  H2O  (p.  155) 
CuO  4-  2  HN03 >-  Cu(N03)2  +  H2O  (p.  129) 

NaOH  +  HC1 >•  NaCl  +  H2O  (p.  109) 

CaC03  +  2  HC1 >-  CaCl2  +  H2O  +  CO2  (p.  198) 

2.  Insoluble  salts.    A  very  large  number  of  salts  are  in- 
soluble in  water,  and  these  can  be  made  by  precipitation. 
All  that  it  is  necessary  to  do  is  to  bring  together  in  solu- 
tion two  salts,  one  of  which  contains  the  desired  metallic 
cation  and  the  other  the  anion.   If,  for  example,  it  is  known 
that  silver  chloride  is  insoluble  in  water,  it  is  to  be  ex- 
pected that  this  salt  will  be  precipitated  on  bringing  to- 
gether in  solution  a  soluble  silver  salt  (AgNO3)   and  a 
soluble  chloride  (NaCl)  : 

AgN03  +  NaCl ^AgCl  +  NaNO3 

If  the  desired  salt  is  also  insoluble  in  strong  acids,  it  may  be 
prepared  by  treating  a  soluble  salt  with  the  proper  acid : 

BaCl2  +  H2S04  — >•  BaS04  +  2  HC1 


274  FIRST  COURSE  IN  CHEMISTRY 

Electrochemical  industries.  To  an  ever-increasing  extent 
electrical  energy  is  being  used  both  in  the  separation  and 
refining  of  the  metals  and  for  the  production  of  many  com- 
pounds. Such  methods  have  in  a  number  of  instances 
already  been  mentioned.  Naturally  these  industries  tend 
to  develop  most  extensively  in  localities  where  water  power 
is  abundant.  Norway  has  many  electrochemical  industries. 
Those  of  the  United  States  and  Canada  center  at  Niagara 
Falls,  the  extensive  power  plants  being  shown  in  Fig.  136. 


FIG.  136.    Electrochemical  power  plants  at  Niagara  Falls 

Solubility  of  salts.  It  will  be  seen  that  a  knowledge  of 
the  solubility  of  various  salts  is  of  great  importance  if  one 
wishes  to  devise  a  means  of  preparing  a  given  salt.  For- 
tunately it  is  possible  to  put  into  brief  form  the  facts  relat- 
ing to  the  solubility  of  the  common  salts,  and  these  rules 
will  have  constant  application  in  the  pages  which  follow. 

1.  Hydroxides.    All  hydroxides  are  insoluble  except  those 
of  ammonium,   sodium,   potassium,  calcium,   barium,   and 
strontium. 

2.  Nitrates.    All  nitrates  are  soluble. 


THE  METALS  275 

3.  Chlorides.    All  chlorides  are  soluble  except  silver  and 
mercurous  chlorides.  (Lead  chloride  is  but  slightly  soluble.) 

4.  Sulfates.    All  sulfates  are  soluble  except  those  of  lead, 
barium,  and  strontium.   (Sulfates  of  silver  and  calcium  are 
only  moderately  soluble.) 

5.  Sulfides.    All  sulfides  are  insoluble  except  those  of  am- 
monium, sodium,  and  potassium.    The  sulfides  of  calcium, 
barium,  strontium,  and  magnesium  are  insoluble  in  water, 
but   are   changed  by  hydrolysis   into  acid  sulfides  which 
are  soluble.    On  this  account  they  cannot  be  prepared  by 
precipitation. 

6.  Carbonates,  phosphates,  and  silicates.    All  normal  car- 
bonates, phosphates,  and  silicates  are  insoluble  except  those 
of  ammonium,  sodium,  and  potassium. 

EXERCISES 

1.  Devise  reactions  for  the  preparation  of  the  following  soluble 
salts :  Ca  (NO3)2,  CuSO4,  ZnCl2,  FeSO4,  and  KNO3. 

2.  Devise  reactions  for  the  preparation  of  the  following  insoluble 
salts :  ri>SO4,  HgCl,  CaCO3,  Fe  (OH)8,  BaSO4. 

3.  How  does  water  power  aid  in  the  electrochemical  industries? 

4.  In  what  processes  already  studied  is  the  electric  current  used? 


CHAPTER  XXIX 
THE  SODIUM  FAMILY 


NAME 

SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POINT 

FIRST  PREPARED 

Lithium 

Li 

6.94 

0.534 

186° 

Arfvedson,  1817 

Sodium 

Na 

23.00 

0.971 

97° 

Davy,  1807 

Potassium 

K 

39.10 

0.862 

62° 

Davy,  1807 

Rubidium 

Rb 

85.45 

1.53 

38° 

Bunsen,  1861 

Caesium 

Cs 

132.81 

1.87 

26° 

Bunsen,  1861 

The  family.  The  metals  appearing  in  the  above  table 
constitute  a  family  in  Group  I  in  the  periodic  arrangement 
of  the  elements.  The  name  alkali  metals  is  often  applied 
to  the  family  for  the  reason  that  the  hydroxides  of  the 
most  familiar  members  of  the  family  (namely,  sodium  and 
potassium)  have  long  been  called  alkalis.  The  alkali  metals 
are  all  readily  acted  upon  both  by  air  and  by  water,  and  are 
not  found  in  a  free  state  in  nature.  Sodium  and  potassium 
are  the  only  important  members  of  the  family. 

SODIUM 

Occurrence.  Large  deposits  of  sodium  chloride  have  been 
found  in  various  parts  of  the  world,  and  the  water  of  the 
ocean  and  of  many  lakes  and  springs  contains  notable 
quantities  of  it.  The  element  also  occurs  as  a  constituent 
of  many  rocks,  and  its  compounds  are  therefore  present  in 
the  soil  formed  by  their  disintegration.  The  nitrate,  the  car- 
bonate, the  borate,  and  many  silicates  also  occur  in  nature. 

276 


THE  SODIUM  FAMILY 


277 


Preparation.    Formerly  sodium  was  prepared  by  reducing 
the  carbonate  with  carbon  : 


NaC0  +  2  C 


2  Na  4-  3  CO 


At  present  it  is  prepared  by  the  electrolysis  of  the  melted 
hydroxide  or  chloride.  Water  must  be  excluded  ;  otherwise 
the  sodium,  as  fast  as  it  is  liberated,  will  react  with  the  water 
to  form  sodium  hydroxide. 

Technical  preparation.  The 
sodium  hydroxide  is  melted 
in  a  cylindrical  iron  vessel  A 
(Fig.  138)  through  the  bottom 
of  which  rises  the  cathode  B. 
The  anodes  C,  several  in  num- 
ber, are  suspended  around  the 
cathode  from  above.  A  cylin- 
drical vessel  E  floats  in  the 
fused  alkali,  directly  over  the 
cathode,  and  under  this  cap 
the  sodium  and  hydrogen  lib- 
erated at  the  cathode  collect. 
The  hydrogen  escapes  by  lift- 
ing the  cover,  and  the  sodium, 
protected  from  the  air  by  the 
hydrogen,  is  from  time  to  time 
skimmed  or  drained  off.  Oxy- 
gen is  set  free  upon  the  anode 
and  escapes  into  the  air  through 
the  opening  F  without  coming 
into  contact  with  the  sodium 
or  the  hydrogen. 


FIG.  137.    Robert  Wilhelm  Bunsen 
(1811-1899) 

A  distinguished  German  chemist  who 
discovered  rubidium  and  caesium,  in- 
vented the  laboratory  burner,  and 
contributed  to  many  advances  in 
chemistry 


Properties.  Sodium  is  a  silver-white  metal  about  as  heavy 
as  water,  and  so  soft  that  it  can  be  molded  easily  by  the 
fingers  or  pressed  into  wire.  It  is  very  active  chemically, 
combining  with  most  of  the  nonmetallic  elements,  such  as 


278 


FIRST  COURSE  IN  CHEMISTRY 


oxygen  and  chlorine,  with  great  energy.  It  will  often  dis- 
place hydrogen  or  metals  combined  with  other  elements 
(p.  112),  and  is  thus  able  to  decompose  water  and  the  ox- 
ides and  chlorides  of  many  metals.  It  forms  many  useful 
compounds,  nearly  all  of  which  are  soluble  in  water. 

Sodium  peroxide  (Na202).  Since  so- 
dium is  a  univalent  element,  we  should 
expect  it  to  form  an  oxide*  of  the  for- 
mula Na2O.  While  such  an  oxide  can 
be  prepared,  the  peroxide,  Naf;O2,  is 
much  better  known.  This  is  a  yel- 
lowish-white powder  made  by  burning 
sodium  in  air.  Its  chief  use  is  as  an 
oxidizing  agent.  When  heated  with 
oxidizable  substances,  it  gives  up  a 
part  of  its  oxygen,  as  shown  in  the 

equation 
FIG.  138.    An  electro- 
lytic cell  for  the  prepa-  Na,,O2 *"  -^a2^  +  [^] 

ration  of  sodium 

When    sodium   peroxide    is    brought 

into  contact  with  water,  oxygen  is  evolved.  On  this  account 
it  is  sometimes  used  in  the  preparation  of  oxygen : 

2  Na2O2  +  2  H2O >•  4  NaOH  +  O2 

Sodium  hydroxide  (caustic  soda)  (NaOH).  Sodium  hydrox- 
ide is  prepared  commercially  by  several  processes. 

1.  In  the  older  process,"  still  in  extensive  use,  sodium 
carbonate  is  treated  with  calcium  hydroxide.  Calcium  car- 
bonate is  precipitated  according  to  the  equation 

Na2CO3  -f  Ca(OH)2 >-  CaCOg  +  2  NaOH 

The  dilute  solution  of  sodium  hydroxide,  filtered  from  the 
calcium  carbonate,  is  evaporated  to  a  paste  and  then  poured 
into  molds  to  solidify.  It  is  sold  in  the  form  of  slender  sticks. 


THE  SODIUM  FAMILY  279 

2.  The  newer  methods  depend  upon  the  electrolysis  of 
sodium  chloride : 

2  NaCl  +  2  H2O >-  2  NaOH  +  H2  +  C12 

The  chief  difficulty  in  this  process  is  to  prevent  the  result- 
ing sodium  hydroxide  and  chlorine  from  coming  in  contact 
and  acting  upon  each  other.  This  difficulty  is  overcome  by 
separating  the  cathode  compartment  (in  which  the  sodium  hy- 
droxide is  produced)  from  the  anode  compartment  (in  which 
the  chlorine  is  evolved)  by  means  of  a  porous  partition. 

Properties.  Sodium  hydroxide  is  a  brittle,  white  crystal- 
line substance  which  rapidly  absorbs  water  and  carbon 
dioxide  from  the  air.  As  the  name  caustic  soda  indicates,  it 
is  a  very  corrosive  substance,  having  a  disintegrating  action 
on  most  animal  and  vegetable  tissues.  In  solution  it  is  a 
strong  base.  It  is  used  in  a  great  many  chemical  industries, 
its  chief  use  being  in  the  manufacture  of  soap  (p.  291). 
As  a  household  article  it  is  sold  under  the  name  of  lye. 

Sodium  chloride  (common  salt)  (NaCl).  Sodium  chloride, 
or  common  salt,  is  very  widely  distributed  in  nature.  Thick 
strata,  evidently  deposited  by  the  evaporation  of  salt  water, 
are  found  in  many  places.  In  the  United  States  the  most 
important  localities  for  salt  are  New  York,  Michigan,  Ohio, 
and  Kansas.  Sometimes  the  salt  is  mined,  especially  if  it 
is  in  the  pure  form  called  rock  salt.  More  frequently  a 
strong  brine  is  pumped  from  deep  wells  sunk  into  the  salt 
deposit.  The  brine  is  evaporated  either  by  heating  or,  in 
the  preparation  of  the  coarser  grades  of  salt,  by  simply  ex- 
posing the  brine  to  the  air  (Fig.  139).  Salt  crystallizes  in 
the  form  of  small  cubes. 

Uses  of  salt.  Since  salt  is  so  abundant  in  nature,  it  is 
used  in  the  preparation  of  nearly  all  compounds  containing 
either  sodium  or  chlorine.  These  include  many  substances 


280  FIRST  COURSE  IK  CHEMISTRY 

of  the  highest  importance  to  civilization,  such  as  soap,  glass, 
hydrochloric  acid,  soda,  and  bleaching  powder.  Enormous 
quantities  of  salt  are  therefore  produced  each  year.  Small 
quantities  are  essential  to  the  life  of  man  and  animals. 
Pure  salt  does  not  absorb  moisture ;  the  fact  that  ordinary 
salt  becomes  moist  when  exposed  to  air  is  not  due  to  a 
property  of  the  salt  but  to  impurities  occurring  in  it,  espe- 
cially to  the  presence  of  calcium  and  magnesium  chlorides. 


FIG.  139.   The  evaporation  of  salt  brine  in  the  open  air  (New  York  State) 

Sodium  sulfate  (Na2S04).  This  salt  is  prepared  by  the 
action  of  sulfuric  acid  upon  sodium  chloride,  hydrogen 
chloride  being  formed  at  the  same  time  (p.  175) : 

2  NaCl  +  H2SO4 >  Na2SO4  +  2  HC1 

The  anhydrous  salt  is  a  white  solid.  It  is  readily  soluble 
in  water,  and  under  ordinary  conditions  crystallizes  out 
as  the  hydrate  Na.2SO4  •  10  H2O  (known  as  Glauber  s  salt). 
Large  quantities  of  sodium  sulfate  are  used  in  making 
sodium  carbonate  and  glass.  The  salt  is  also  used  in 
medicine. 

Sodium  carbonate  (Na2C03).  There  are  two  different 
methods  now  employed  in  the  manufacture  of  this  salt. 


THE  SODIUM  FAMILY  281 

1.  Leblanc  process.    This  older  process  involves  several 
distinct  reactions,  as  shown  in  the  following  equations  : 

a.  Sodium  chloride  is  first  converted  into  sodium  sul- 

fofp  • 

2  NaCl  +  H2S04  --  >•  Na2SO4  +  2  HC1 

b.  The  sodium   sulfate   is  next  reduced  to  sulfide  by 
heating  it  with  carbon  : 

Na2S04  +  2  C  —  >•  Na2S  +  2  CO2 

c.  The  sodium  sulfide  is  then  heated  with  calcium  car- 
bonate, when  the  following  reaction  takes  place  : 

Na2S  +  CaC03  -  *  CaS  +  Na2CO3 

2.  Solvay  process.    This  more  modern  process  depends 
upon  reactions  taking  place  in  solution,  and  represented  in 
the  equations 

NaCl  +  NH4HCO3  —  >-  NaHCO3  +  NH4C1  (1) 

2  NaHC03  -  >•  Na2C03  +  H2O  +  CO2         (2) 

When  concentrated  solutions  of  sodium  chloride  and  of 
ammonium  hydrogen  carbonate  are  brought  together,  the 
sparingly  soluble  sodium  hydrogen  carbonate  is  precipi- 
tated as  represented  in  equation  (1).  This,  by  heating,  is 
converted  into  the  normal  carbonate  as  indicated  in  equa- 
tion (2).  The  ammonium  chloride  formed  (equation  (1)) 
is  treated  with  lime  (p.  122),  ammonia  being  liberated. 
This  ammonia,  together  with  water  and  the  carbon  dioxide 
generated  as  indicated  in  equation  (2),  combine  to  form 
ammonium  hydrogen  carbonate, 


NH3+C0 

This  is  treated  with  salt,  and  the  process  is  begun  over  again. 


282 


FIRST  COURSE  IN  CHEMISTRY 


Historical.  In  former  times  sodium  carbonate  was  made  by 
burning  seaweeds  and  extracting  the  carbonate  from  their  ash. 
On  this  account  the  salt  was  called  soda  ash,  and  the  name  is 

still  in  common  use.    During  the 
French   Revolution   this    supply 


was    cut   off,   and    in    behalf   of 


the  French  government  Leblanc 
(Fig.  140)  made  a  study  of  meth- 
ods of  preparing  the  carbonate 
directly  from  salt.  As  a  result 
he  devised  the  method  which 
bears  his  name,  and  which  was 
used  exclusively  for  many  years. 
It  has  been  replaced  to  a  large 
extent  by  the  Solvay  process, 
which  was  devised  by  the  Belgian 
chemist  Solvay  in  1863. 

By-products.  The  substances 
obtained  in  a  given  process,  aside 
from  the  main  product,  are  called 
the  ly-products.  The  success  of 
many  processes  depends  upon 
the  value  of  the  by-products 
formed.  Thus,  hydrochloric  acid, 
a  by-product  in  the  Leblanc  proc- 
ess, is  valuable  enough  to  make 
the  process  pay,  even  though  so- 
dium carbonate  can  be  made  more 
cheaply  in  other  ways. 

Properties  of  sodium  carbon- 
ate. Sodium  carbonate  forms 
large  crystals  of  the  formula 
Na2CO3  •  10  H2O,  known  as 

washing  soda,  or  sal  soda.  Its  solution  in  water  has  a  mild 
alkaline  reaction,  and  is  used  for  laundry  purposes.  Mere 
mention  of  the  fact  that  it  is  used  in  the  manufacture  of 


FIG.  140.    Nicolas  Leblanc 
(1742-1806) 

Inventor  of  the  first  method  of 

preparing  sodium  carbonate  from 

salt.    (Statue  erected  in  Paris) 


THE  SODIUM  FAMILY  283 

glass,  soap,  and  many  chemical  reagents  will  indicate  its 
importance  in  the  industries.  It  is  one  of  the  few  soluble 
carbonates. 

Sodium  hydrogen  carbonate  (NaHC03).  This  salt,  called 
bicarbonate  of  soda,  or  hiking  soda,  is  made  by  the  Solvay 
process,  as  explained  above,  or  by  passing  carbon  dioxide 
into  concentrated  solutions  of  sodium  carbonate: 

Na2C08  +  H20  +  C()2 *  2  NaHCO3 

It  is  an  essential  constituent  of  all  baking  powders  (p.  330). 


FIG.  141.    A  sodium  nitrate  deposit  in  Chile 

Sodium  nitrate  (NaN03).  This  substance,  known  also  as 
Chile  saltpeter,  is  found  in  nature  in  certain  arid  regions, 
where  apparently  it  has  been  formed  by  the  decay  of  or- 
ganic substances  in  the  presence  of  air  and  sodium  salts. 
The  largest  deposits  are  in  Chile,  and  most  of  the  nitrate 
of  commerce  comes  from  that  country.  Fig.  141  shows  a 
deposit  of  sodium  nitrate  in  Chile  after  it  has  been  broken 
apart  by  explosives.  Smaller  deposits  occur  in  California 
and  Nevada.  The  commercial  salt  is  prepared  by  dissolving 


284  FIKST  COUKSE  IN  CHEMISTRY 

the  crude  nitrate  (known  as  caliche)  in  water,  allowing  the 
insoluble  earthy  materials  to  settle,  and  evaporating  to 
crystallization  the  clear  solution  so  obtained.  The  soluble 
impurities  remain  for  the  most  part  in  the  mother  liquors. 

Since  this  salt  is  the  only  nitrate  found  extensively  in 
nature,  it  is  the  material  from  which  other  nitrates,  as  well 
as  nitric  acid,  are  prepared.  Enormous  quantities  are  used 
as  a  fertilizer  and  in  the  manufacture  of  sulfuric  acid. 

Sodium  thiosulfate  (Na2S203).  This  salt  is  made  by  boil- 
ing a  solution  of  sodium  sulfite  with  sulfur : 

Na2S03  +  S *NagS808 

The  hydrate  Na2S2O3  •  5  H2O  is  frequently  called  sodium 
hyposulfite,  or  simply  hypo.  It  is  used  in  photography 
(p.  367),  and  in  the  bleaching  industry  to  absorb  the  ex- 
cess of  chlorine  which  is  left  upon  the  bleached  fabrics. 

Other  compounds  of  sodium.  (1)  Sodium  sulfite  (Na2SO3). 
(2)  Normal  sodium  phosphate  (NagPO4).  This  salt  readily 
hydrolyzes,  forming  a  basic  solution.  (3)  Disodium  phos- 
phate (Na0HPO4).  This  is  the  most  common  of  the  phos- 
phates of  sodium,  and  is  usually  known  simply  as  sodium 
phosphate. 

POTASSIUM 

Occurrence.  Potassium  is  a  rather  abundant  element, 
being  a  constituent  of  many  igneous  rocks,  such  as  the  feld- 
spars and  micas.  Very  large  deposits  of  the  chloride  and 
the  sulfate,  associated  with  compounds  of  calcium  and  mag- 
nesium, are  found  at  Stassfurt,  Germany,  and  are  known 
as  the  Stassfurt  salts. 

The  natural  decomposition  of  rocks  containing  potassium 
gives  rise  to  various  compounds  of  the  element  in  all  fertile 
soils.  Its  soluble  compounds  are  absorbed  by  growing 


THE  SODIUM  FAMILY 


285 


plants  and  built  up  into  complex  vegetable  tissues ;  when 
these  are  burned,  the  potassium  remains  in  the  ash  in  the 
form  of  carbonate.  The  crude  carbonate  can  be  separated 
from  wood  ashes  by  dissolving  it  in  water.  This  was  for- 
merly the  chief  source  of  potassium  compounds,  but  they  are 
now  prepared  mostly  from  the  salts  of  the  Stassfurt  deposits. 


FIG.  142.    Mining  Stassfurt  salts  for  use  in  the  manufacture  of  fertilizers 

Potassium  in  sea  plants.  While  sodium  rather  than  potassium 
is  likely  to  be  present  in  sea  plants,  nevertheless  some  of  these 
plants,  such  as  the  giant  algae  of  the  California  coast,  contain 
potassium  chloride,  amounting  in  some  cases  to  30  per  cent  of 
their  dry  weight.  It  is  thought  that  perhaps  these  may  consti- 
tute a  commercial  source  of  potassium  compounds,  large  quan- 
tities of  which  are  used  in  fertilizers.  At  present,  however,  the 
Stassfurt  salts  constitute  by  far  the  cheapest  source,  although 
the  dried  algae  are  being  used  to  some  extent  as  a  fertilizer. 


286  FIRST  COURSE  IN  CHEMISTRY 

Stassfurt  salts.  These  salts  form  very  extensive  deposits  in 
middle  and  north  Germany,  the  most  noted  locality  for  working 
them  being  at  Stassfurt.  The  deposits  are  very  thick  and  rest 
upon  an  enormous  layer  of  common  salt.  They  are  in  the  form 
of  a  series  of  strata,  each  consisting  largely  of  a  single  mineral 
salt.  While  these  strata  are  salts  from  a  chemical  standpoint, 
they  are  as  solid  and  hard  as  many  kinds  of  stone  and  are  mined 
as  stone 'or  coal  would  be  (Fig.  142).  The  chief  minerals  of 
commercial  importance  in  these  deposits  are  the  following : 

Sylyite KC1 

Kainite KC1  •  MgSO4  •  3  H2O 

Carnallite KC1  •  MgCl2  •  6  II2O 

Kieserite MgSO4  •  H2O 

Preparation  and  properties.  Potassium  is  prepared  by 
methods  similar  to  those  used  in  the  preparation  of  sodium. 
It  is  more  active  than  sodium ;  otherwise  the  properties  of 
the  two  metals  are  alike. 

Compounds.  The  compounds  of  potassium  are  so  similar 
in  properties  to  the  corresponding  compounds  of  sodium  that 
for  many  purposes  for  which  they  are  used  they  can  be 
interchanged.  The  compounds  of  potassium,  being  as  a  rule 
more  expensive,  are  not  so  widely  used  as  those  of  sodium. 

Potassium  hydroxide  (caustic  potash)  (KOH).  Potassium 
hydroxide  is  prepared  by  methods  exactly  similar  to  those 
used  in  the  preparation  of  sodium  hydroxide,  which  com- 
pound it  closely  resembles  hi  both  physical  and  chemical 
properties.  It  is  not  used  to  any  very  great  extent,  being 
replaced  by  the  cheaper  sodium  hydroxide. 

Action  of  the  halogen  elements  on  bases.  When  any  one  of 
the  three  halogen  elements  —  chlorine,  bromine,  or  iodine  —  is 
added  to  a  solution  of  a  base  such  as  the  hydroxide  of  sodium, 
of  potassium,  or  of  calcium,  a  reaction  takes  place,  the  nature  of 
which  depends  upon  the  temperature.  Thus,  when  chlorine  is 


THE  SODIUM  FAMILY  287 

passed  into  a  cold  solution  of  potassium  hydroxide,  the  reac- 
tion expressed  by  the  following  equation  takes  place : 

2  KOH  +  H, >•  KC1  +  KC10  +  H2O  (1) 

If  the  solution  of  hydroxide  is  hot,  the  potassium  hypochlorite 
(K(110)  formed  according  to  equation  (1)  breaks  down  as  fast 
as  formed :  g  ^  _^  ^^  +  ^  ^  (2) 

By  means  of  these  reactions  one  can  prepare  the  chloride,  hypo- 
chlorite, and  chlorate  of  sodium,  of  potassium,  and  of  calcium. 
By  substituting  bromine  or  iodine  for  chlorine  the  correspond- 
ing compounds  of  those  elements  are  obtained.  Some  of  these 
compounds  can  be  obtained  in  cheaper  ways. 

Potassium  chlorate  (KC103).  This  salt,  as  has  just  been 
explained,  can  be  made  by  the  action  of  chlorine  on  concen- 
trated solutions  of  potassium  hydroxide.  It  is  a  white  crys- 
talline substance,  and  is  used  chiefly  as  an  oxidizing  agent 
in  the  manufacture  of  matches,  fireworks,  and  explosives ; 
it  is  also  used  in  the  preparation  of  oxygen  and  in  medicine. 

Potassium  nitrate  (saltpeter)  (KN03).  This  salt  is  found 
native  in  some  regions  where  the  climate  is  hot  and  dry, 
being  formed  by  the  decay  of  nitrogenous  organic  matter 
in  the  presence  of  earthy  material  containing  potassium. 
The  saltpeter  used  in  making  gunpowder  was  formerly 
made  by  imitating  these  conditions.  At  present  it  is  pre- 
pared by  the  action  of  sodium  nitrate  upon  potassium  chlo- 
ride (the  former  compound  being  obtained  from  Chile  and 
the  latter  from  the  Stassfurt  deposits)  : 

NaNO3  4-  KC1 >-  KNO3  +  NaCl 

Potassium  nitrate  is  a  white  solid.  It  is  an  excellent 
oxidizing  agent,  and  its  chief  use  is  in  the  manufacture  of 
gunpowder  (p.  294).  For  this  purpose  it  is  preferable  to 
sodium  nitrate,  since  the  latter  is  slightly  deliquescent 


288  FIRST  COURSE  IN  CHEMISTRY 

and  powder  made  from  it,  if  exposed  to  air,  soon  becomes 
moist  and  unfit  for  use.  Small  amounts  of  the  nitrate  are 
also  used  in  medicine  and  as  a  preservative  for  meats, 
especially  for  corned  beef. 

Other  compounds  of  potassium.  Potassium  chloride  (KC1), 
potassium  bromide  (KBr),  and  potassium  iodide  (KI)  are  all 
white  solids.  The  bromide  and  the  iodide  are  both  used  in  the 
preparation  of  photographic  reagents  and  in  medicine.  Potas- 
sium carbonate  (K2C03),  potassium  bicarbonate  (KHC03),  potas- 
sium sulfate  (K2S04),  and  potassium  bisulfate  (KHS04j  are  all 
well-known  compounds.  They  are  white  solids,  readily  soluble 
in  water. 

THE  AMMONIUM  COMPOUNDS 

Composition.  As  explained  in  a  previous  chapter,  when 
ammonia  is  passed  into  water  the  two  combine  to  form  the 
base  NH4OH,  known  as  ammonium  hydroxide.  When  this 
base  is  neutralized  with  acids,  there  are  formed  the  corre- 
sponding salts,  known  as  the  ammonium  salts.  Since  the 
ammonium  radical  is  univalent,  ammonium  salts  resemble 
those  of  the  alkali  metals  in  their  formulas  ;  they  also  re- 
semble the  latter  salts  very  much  in  their  chemical  proper- 
ties, and  may  be  conveniently  described  in  connection  with 
them.  Among  the  ammonium  salts  the  chloride,  sulfate, 
carbonate,  and  sulfide  are  the  most  familiar. 

Ammonium  chloride  (sal  ammoniac)  (NH4C1).  This  is  a 
white  solid.  When  heated  it  partly  decomposes  into  ammo- 
nia and  hydrogen  chloride,  which  recombine,  as  the  tern- 
perature  falls  :  NH  ^ 

~ 


This  salt  is  used  in  soldering,  since  the  hydrogen  chloride 
evolved  by  the  heat  removes  the  oxide  from  the  surface  of 
the  metals.  It  is  also  used  in  making  dry  cells,  in  medicine, 
and  as  a  chemical  reagent. 


THE  SODIUM  FAMILY 


289 


Ammonium  sulfate  ((NHJ2S04).  This  salt  is  prepared  at 
low  cost,  and  is  especially  valuable  as  a  fertilizer. 

The  carbonates  of  ammonium.  Both  the  normal  carbonate 
(NH4)2CO3  and  the  acid  carbonate  NH4HCO3  are  white 
solids,  readily  soluble  in  water.  The  normal  carbonate 
slowly  decomposes  into  the  acid  carbonate,  evolving 

ammonia : 

(NH4),CO.  — *•  NH.HCO.  +  NH. 

Both  carbonates  are  used  as  chemical  reagents. 

Ammonium  sulfides.  When  hydrogen  sulfide  is  passed 
into  aqua  ammonia,  a  solution  containing  ammonium  sul- 
fide ((NH4).,S)  and  ammonium  acid  sulfide  (NH4HS)  is 

obtained : 

NH4OH  +  H0S H  NH4HS  +  H2O 

2  NH4OH  +  H2S >-  (NH4)2S  +  2  H2O 

This  solution  is  usually  known  simply  as  ammonium  sulfide, 
and  is  used  as  a  reagent  in  testing  for  certain  metals.  When 
exposed  to  the  air  it  slowly 
decomposes,  and  the  sulfur 
liberated  in  the  process 
combines  with  the  com- 
pounds present,  forming 
different  sulfides,  such  as 
(NH4)2S2and(NH4).2S3,or, 
in  general,  (NH4)2SX.  The 
resulting  solution  is  yellow 
and  is  termed  yellow  am- 
monium sulfide  or  ammo- 
nium polysulfide.  FIG.  143.  Method  of  making  a  flame  test 

Flame  reactions.  A  number  of  the  metals,  when  volatil- 
ized in  a  colorless  flame,  such  as  that  of  a  Bunsen  burner, 
impart  a  characteristic  color  to  the  flame.  Thus,  sodium 
(or  any  of  its  compounds  that  will  volatilize  in  the  heat  of 


290  FIRST  COUESE  IN  CHEMISTRY 

the  flame)  imparts  to  tlie  flame  a  strong  yellow  color.  Po- 
tassium and  its  compounds  color  the  flame  a  pale  violet, 
and  lithium  colors  it  a  deep  crimson  red. 

Advantage  is  taken  of  these  facts  in  testing  for  the  pres- 
ence of  the  elements  in  different  substances.  The  test  is 
best  made  by  using  a  platinum  wire,  one  end  of  which  is 
fused  into  a  piece  of  glass  tubing  that  serves  as  a  handle. 
The  other  end  of  the  wire  is  dipped  into  water  and  rubbed 
in  the  substance  to  be  tested  (or  dipped  into  a  concen- 
trated solution  of  the  substance),  and  the  wire  with  the 
adhering  particles  is  held  in  the  outer  edge  of  the  base 
of  the  Bun  sen  flame  (Fig.  143). 

EXERCISES 

1.  What  is  an  alkali?    What  does  the  word  alkali,  signify  (see 
dictionary)?    Can  a  metal  itself  be  an  alkali? 

2.  AVhat  carbonates  are  soluble? 

3.  Account  for  the  fact  that  solutions  of  sodium  carbonate  and 
potassium  carbonate  are  basic. 

4.  What  nomiietallic  element  is  obtained  from  the  deposits  of 
Chile  saltpeter? 

5.  What  would  take  place  if  a  bit  of  potassium  hydroxide  were 
left  exposed  to  the  air  ? 

6.  What  substances  already  studied  are  prepared  from  the  fol- 
lowing compounds  :   ammonium  chloride  ;   ammonium  nitrate ;   am- 
monium nitrite  ;   sodium  nitrate  ? 

7.  What  weight  of  sal  soda  can  be  prepared  from  1  ton  of  salt? 
,4ns.4894.9  Ib.  What  weight  of  it  is  water  of  hydration?  .4ns. 3081.8  Ib. 

TOPICS  FOR   TlIKMKS 

The  kelp  industry  (Koscoe  and  Scliorlemmer,  Chemistry;  see  also 
encyclopedia). 

The  Stassfurt  deposits  (Roscoe  and  Scliorlemmer,  Chemistry;  see 
also  encyclopedia). 

The  salt  wells  of  the  United  States.  (Write  to  your  state  geologist 
for  sources  of  information.) 


CHAPTER  XXX 
SOAP;  GLYCERIN;  EXPLOSIVES 

Composition  of  soap,  and  materials  used  in  its  preparation. 
Soap  is  a  mixture  of  the  sodium  and  potassium  salts  of 
oleic,  palmitic,  and  stearic  acids.  The  essential  materials 
used  in  the  preparation  of  soap  are  as  follows: 

1.  Fat  or  oil.    As  shown  on  page  237,  fats  and  oils  are 
largely  mixtures  of  olein,  palmitin,  and  stearin.  The  cheaper 
grades  of  these  are  used  in  making  soap.    Those  commonly 
employed  are  a  low  grade  of  animal  fat  (tallow)  and  the 
cheaper  vegetable  oils,  such  as  cottonseed  oil,  coconut  oil, 
and  palm  oil. 

A  low-grade  fat  obtained  from  garbage  is  also  used  in  making 
soap  and  candles.  When  the  garbage  is  heated  with  water,  the 
fat  rises  to  the  top  and  is  skimmed  off.  The  remaining  matter 
is  used  as  a  constituent  of  fertilizers. 

2.  Alkali.    The  alkali  used  is  the  hydroxide  of   either 
sodium  or  potassium.    Sodium  hydroxide  is  nearly  always 
used,  since  it  gives  a  hard  soap,  while  potassium  hydroxide 
gives  a  soft  soap. 

Reaction  taking  place  in  the  preparation  of  soap.  When 
the  fat  and  alkali  are  heated  together  under  proper  condi- 
tions, the  olein,  palmitin,  and  stearin  present  in  the  fat  are 
decomposed,  forming  glycerin,  together  with  sodium  oleate, 
sodium  palmitate,  and  sodium  stearate.  A  mixture  of  these 
three  salts  constitutes  soap.  The  reactions  may  be  illustrated 

291 


292 


FIRST  COURSE  IN  CHEMISTRY 


by  the  following  equation,  which   represents  the  change 
taking  place  when  stearin  is  heated  with  sodium  hydroxide  : 

C3H5(CI8H8502)3  +  3  NaOH  -+C.II.(OH),+  3 


stearin 


sodium  hydroxide 


glycerin 


sodium  stearate 


In  this  reaction  the  fat  is  said  to  be  saponified,  and  the 
process  is  known  as  saponification. 

Commercial  manufacture  of  soap.  The  oil  or  melted  fat  is 
poured  into  large  iron  kettles  together  with  a  solution  of  so- 
dium hydroxide  containing  about  one  fourth  of  the  amount  of 

alkali  necessary  to  saponify  the  fat. 
As  a  rule,  the  kettles  are  very  large 
(Fig.  144),  500,000  Ib.  or  more  of 
soap  being  made  in  some  of  them  in 
a  single  heating.  They  are  provided 
with  coils  of  steam  pipe  for  heating 
the  mixture.  The  fat  and  alkali  are 
stirred  by  forcing  air  or  live  steam 
into  the  bottom  of  the  mixture.  As 
the  heating  continues,  the  remainder 
of  the  alkali  is  added.  The  reaction 
is  complete  in  from  two  to  five  days. 
The  soap  is  next  removed,  or 
salted  out,  from  the  mixture.  This 
process  consists  in  adding  salt  and 
again  heating.  After  a  time  the 
soap  rises  to  the  top  of  the  liquid 
(or  spent  lye,  as  it  is  called).  The 
soap  so  obtained  is  purified  and 
then  run  into  a  mixing  machine 
(crutcher).  Here  it  is  mixed  with 

any  appropriate  material  which  it  is  desired  to  add,  such  as 
perfume,  borax,  sodium  silicate,  or  sodium  carbonate.  It  is 
then  run  into  large  molds  to  harden,  after  which  it  is  cut  and 
pressed  into  cakes  of  the  desired  size.  The  glycerin  formed 
in  the  reaction  is  separated  from  the  spent  lye  by  distillation. 


FIG.  144.    A  kettle  used  in 
making  soap 


SOAP;  GLYCERIN;  EXPLOSIVES  293 

Varieties  of  soap.  Transparent  soaps  are  ordinarily  made  by 
dissolving  soap  in.  alcohol.  The  solution  is  filtered  and  the 
excess  of  alcohol  removed  by  distillation.  Castile  soaps  are 
made  from  mixtures  of  olive  oil  with  cheaper  oils.  The  color 
of  mottled  soaps  is  produced  by  the  addition  of  ferrous  sulfate, 
Prussian  blue,  or  some  similar  pigment.  Floating  soaps  owe 
their  lightness  to  bubbles  of  air.  Naphtha  soaps  contain  from 
5  per  cent  to  10  per  cent  of  petroleum  naphtha.  Scouring  soaps 
contain  from  5  per  cent  to  10  per  cent  of  soap  and  from  80  per 
cent  to  90  per  cent  of  some  abrasive  material  such  as  fine  sand 
or  volcanic  ash.  Sometimes  a  small  percentage  of  sodium  car- 
bonate is  also  present.  Soap  powders  are,  as  a  rule,  sodium 
carbonate  or  a  mixture  of  sodium  carbonate  and  ground  soap. 

Properties  of  soap.  Soap  dissolves  in  soft  waters,  giving 
a  slightly  alkaline  solution  due  to  hydrolysis.  If  an  acid, 
such  as  hydrochloric  acid,  is  added  .to  the  aqueous  solu- 
tion, the  organic  acids  are  liberated  from  their  salts  and 
are  precipitated  in  the  form  of  white  insoluble  solids  : 


.0,  +  HC1  —  »-  NaCl  +  H  -C]8H85O2 

sodium  stearate  stearic  acid 

The  calcium  and  magnesium  salts  of  oleic,  palmitic,  and 
stearic  acids  are  insoluble  in  water,  and  are  therefore  pre- 
cipitated when  a  calcium  or  magnesium  compound  is  added 
to  an  aqueous  solution  of  soap  : 

2  NaC^H.O,  +  CaCl,  —  *  2  NaCl  +  CaCC^O,), 

sodium  stearate  calcium  stearate 

It  is  due  to  this  fact  that  soaps  do  not  lather  with  hard 
waters  but  form  a  curdy  precipitate,  since  such  waters 
always  contain  calcium  and  magnesium  salts  in  solution. 

Cleansing  action  of  soap.  Attention  has  been  called  to 
the  property  possessed  by  soap  of  aiding  in  the  formation 
of  emulsions  (Fig.  50).  The  cleansing  power  of  soap  is 


294  FIRST  COUKSE  IN  CHEMISTRY 

largely  due  to  this  fact.  When  soap  is  rubbed  upon  the 
skin,  any  oily  substances  present  are  emulsified  by  the  soap 
solution  and  washed  away. 

Candles.  When  fats  are  heated  with  steam,  they  are  de- 
composed into  glycerin  and  free  acid  as  follows : 

C.H.(CUH,,0,).  +  3  H,0  — H C.H.(OH).  +  3  H  .C18H85O2 

stearin  glycerin  stearic  acid 

The  solid  fatty  acids  thus  obtained,  mixed  with  paraffin, 
are  used  in  making  candles.  The  glycerin  is  easily  sepa- 
rated, and  this  process  and  the  process  of  soap-making 
constitute  the  commercial  sources  of  this  compound. 

Glycerin  (C3H5(OH)3).  This  is  a  colorless  oily  liquid 
having  a  sweet  taste.  Nitric  acid  reacts  with  it  just  as 
with  a  base,  forming  the  nitrate  C8H5(NO8)8: 

C.H.(OH),  +  3  HNO,  — v  C,H.(SO,)t  +  8  H,O 

The  nitric  acid  used  in  this  reaction  is  mixed  with  a  little 
sulfuric  acid,  the  latter  serving  to  absorb  the  water  gener- 
ated. Glycerin  nitrate  is  a  slightly  yellowish  oil  commonly 
known  as  nitroglycerin.  It  is  very  explosive  and  is  used  in 
the  manufacture  of  dynamite  (p.  296). 

Explosives.  An  explosion  is  caused  by  a  very  rapid  chem- 
ical reaction  which  results  in  the  formation  of  large  volumes 
of  gases  from  liquid  and  solid  materials  called  explosives. 
The  greater  the  volume  change  and  the  more  rapidly  it  is 
produced,  the  more  violent  the  explosion.  The  most  com- 
mon of  the  manufactured  explosives  are  the  following : 

1.  Gunpowder.  Ordinary  gunpowder  is  an  intimate  mixture 
of  potassium  nitrate,  sulfur,  and  charcoal.  When  ignited,  com- 
plicated reactions  occur,  the  principal  change  being  indicated 
in  the  following  equation  : 

2  KN0  +  3  C  +  S >•  KS  +  3  CO  +  N2 


SOAP;  GLYCERIN;  EXPLOSIVES  295 

The  explosive  effects  are  due  to  the  sudden  liberation  of  the 
highly  heated  carbon  dioxide  and  nitrogen.  The  presence  of 
the  solid  particles  of  the  sulfide  of  potassium  causes  the  smoke 
which  accompanies  the  explosion. 

2.  Nitrocellulose.  The  nitrocellulose  usually  used  in  explo- 
sives has  the  formula  C6H70./]ST03)3.  This  substance  (p.  225) 
is  a  far  more  powerful  explosive  than  gunpowder.  If  ignited,  it 
will,  under  ordinary  conditions,  burn  quietly.  If  subjected  to 
a  sudden  shock  (such  as  may  be  produced  by  the  explosion  of  a 
small  percussion  primer),  the  nitrocellulose  decomposes  with 
enormous  violence.  The  products  of  the  decomposition  are 


FIG.  145.    Powder  grains  for  large  guns  (natural  size) 

all  colorless  gases ;  hence  the  use  of  this  explosive  in  making 
smokeless  gunpowder.  When  used  for.  this  purpose,  it  is  neces- 
sary to  modify  the  pure  material  somewhat,  as  otherwise  the 
violence  of  the  explosion  would  shatter  any  firearms  in  which 
the  powder  was  used.  This  is  done  by  mixing  nitrocellulose 
with  sufficient  alcohol  and  ether  to  form  a  plastic  mass.  This 
is  then  molded  into  the  form  of  rods  (grains)  with  a  number  of 
perforations  through  the  rods.  Fig.  145  shows  the  form  of  the 
grains  used  in  the  large  guns  of  our  navy. 

3.  Nitroglycerin  (C3H5(N03)3).  This  resembles  nitrocellulose 
in  the  violence  of  its  explosive  effects.  The  changes  taking  place 
in  its  decomposition  are  represented  in  a  general  way  in  the 
following  equation  : 

4  C3H5(N03)3 >- 12  C02  +  6  N2  +  10  H20  +  02 


296  FIRST  COURSE  IN  CHEMISTRY 

One  volume  of  nitrogiycerin  yields  on  explosion  about  1300 
volumes  of  gas,  which  is  expanded  by  the  heat  of  the  reaction 
to  over  10,000  volumes.  Pure  nitrogiycerin  is  very  dangerous 
because  of  the  ease  with  which  it  is  set  off.  Large  quantities 
are  used  in  making  dynamite,  in  which  form  it  is  not  exploded 
so  readily  by  jarring  and  can  be  transported  with  less  danger. 
Ordinary  dynamite  consists  of  a  mixture  of  sodium  nitrate, 
nitrogiycerin,  and  wood  pulp,  the  latter  acting  as  an  absorb- 
ent for  the  nitrogiycerin.  Gelatin  dynamite  consists  of  nitrocel- 
lulose and  nitrogiycerin  mixed  together  to  form  a  jelly  like 
mass.  It  is  a  very  powerful  explosive,  since  it  contains  no  inert 
material. 

4.  Trinitrotoluene  (C7H5(N02)3).  Recently  the  Germans  have 
developed  a  powerful  explosive  known  as  trinitrotoluene.  It  is 
prepared  by  the  action  of  nitric  acid  on  toluene  (p.  232).  It 
is  solid  and  can  be  transported  with  safety. 

EXERCISES 
1.    In  what  way  does  aqua  ammonia  assist  in  the  removal  of 


2.  For  what  is  lye  vised  as  a  household  article  ? 

3.  What  effect  will  the  softening  of  a  city  water-supply  have  on 
soap  consumption? 

4.  Why  will  gas  which  burns  quietly  in  a  stove  explode  violently 
if  a  sufficient  quantity  of  it  is  allowed  to   escape  into  a  room  and 
is  then  ignited  ? 

^5.    Give  the  steps  in  preparing  nitrogiycerin  from  garbage. 

6.  Why  not  use  sodium  nitrate  in  making  gunpowder? 

7.  WThat  are  the  approximate  proportions  in  which  the  constitu- 
ents of  gunpowder  are  mixed  in  its  preparation  (see  equation  for 
reaction  of  explosion)? 

8.  Why  are  some  gunpowders  smokeless  and  others  not? 

TOPICS  FOR  THEMES 

Soap-making  (Rogers  and  Aubert,  Industrial  Chemistry). 
Modern  explosives  (Bird,  Modern  Science  Reader). 


CHAPTER  XXXI 
THE  CALCIUM  FAMILY 


j 

o  ^ 

' 

MILLIGRAMS  SOLU- 

NAME 

c 
A 
fl 

11 

gw 

H 
fe 

BLE  IX  ILITEH  OF 

WATER  AT  18° 

CARBONATE 
DECOMPOSES 

02 

•<£ 

Q 

Sulfate 

Hydroxide 

Calcium 

Ca 

40.07 

1.55 

2070. 

1,670 

At  dull-red  heat 

Strontium 

Sr 

87.63 

2.54 

170. 

7,460 

At  white  heat 

Barium 

Ba 

137.37 

3.75 

2.29 

•  36,300 

Scarcely  at  all 

The  family.-  The  calcium  family  consists  of  the  very 
abundant  metal  calcium  and  the  rarer  metals,  strontium 
and  barium.  Like  the  alkali  metals,  they  are  acted  upon 
by  both  water  and  air,  and  on  this  account  do  not  occur  in 
a  free  state  in  nature.  They  are  bivalent,  so  that  the 
formulas  of  their  salts  differ  from  the  formulas  of  the 
corresponding  salts  of  the  alkali  metals ;  moreover,  their 
normal  carbonates  and  phosphates  are  insoluble  in  water. 


CALCIUM 

Occurrence.  Calcium  is  one  of  the  abundant  elements.  In 
the  form  of  a  carbonate  it  occurs  in  a  number  of  different 
forms,  such  as  limestone  and  marble.  The  most  important 
of  its  mineral  compounds  are  the  following : 


Calcite  .  . 
Phosphorite 
Fluorite 


CaCO3  Wollastonite 

Ca3(PO4)2      Gypsum    .     . 
CaF2  Anhydrite     . 

297 


CaSiO3 
CaS04 .  2  H20 
CaSO, 


298 


FIRST  COURSE  IN  CHEMISTRY 


Preparation.  Calcium  is  prepared  commercially  by  the 
electrolysis  of  the  melted  chloride  in  the  following  way: 

The  apparatus  consists  of  a  cylindrical  iron  vessel  (Fig.  146), 
through  the  bottom  of  which  extends  the  iron  cathode  A.  The 
anodes  B,  B,  several  in  number,  are  placed  about  the  sides  of 
the  vessel.  The  calcium  separates  in  a  molten  condition  at  the 

cathode  .1  ,  and  rises  in  the  form 
of  globules  to  the  lower  surface 
of  a  solid  stick  of  calcium  D, 
suspended  above  the  cathode. 
There  it  is  chilled  by  a  water- 
cooling  device  C,  C,  and  adheres 
to  the  stick  of  calcium,  which 
is  slowly  raised  as  it  increases 
in  length. 

Properties.  Calcium  is  a 
silver-white  metal  which  is 
only  a  little  heavier  than 
water.  It  is  qnite  hard  and 
melts  at  810°.  It  combines 

readily   with   many   of    the 

-,  i      i         •       v    i 

elements,  and  when  ignited, 

burns  in  oxygen  with  daz- 
zling brilliancy.  Like  sodium  and  potassium,  it  decomposes 
water,  forming  a  hydroxide  and  hydrogen.  As  yet  it  has 
few  commercial  applications. 

Calcium  carbonate  (CaC03).  Enormous  quantities  of  cal- 
cium carbonate  occur  in  nature.  Limestone  is  the  most 
abundant  form,  and  sometimes  constitutes  whole  moun- 
tain ranges.  Limestone  is  never  pure  calcium  carbonate, 
but  contains  variable  percentages  of  magnesium  carbon- 
ate, clay,  silica,  and  compounds  of  iron.  Pearls,  coral,  and 
various  shells  are  largely  calcium  carbonate.  Calcite  is 
a  very  pure,  crystalline  form,  and  often  is  found  in  large 


FIG.  146.   The  electrolytic  prepara- 
tion  of  calcium 


THE  CALCIUM  FAMILY 


299 


transparent  crystals  (Fig.  147)  called  Iceland  spar.  Marble 
is  composed  of  very  small  calcite  crystals. 

In  the  laboratory  pure  calcium  carbonate  can  be  prepared 
by  treating  a  soluble  calcium  salt  with  a  soluble  carbonate : 

Na2C03  +  CaCl2 >•  CaCO3  +  2  NaCl 

When  prepared  in  this  way,  it  is  a  soft  white  powder  often 
called  precipitated  chalk,  and  is  much  used  as  a  polishing 
powder  (tooth  powder). 

The  natural  varieties  of 
calcium  carbonate  find  many 
uses,  as  in  the  preparation  of 
lime  and  of  carbon  dioxide; 
in  metallurgical  operations, 
especially  in  blast  furnaces  ; 
in  the  manufacture  of  soda 
and  glass ;  for  building-stone 
and  for  ballast  for  roads. 

Calcium  acid  carbonate 
(Ca(HC03)2).  Calcium  car- 
bonate is  almost  insoluble  in  pure  water.  It  readily  dissolves, 
however,  in  water  which  holds  carbon  dioxide  in  solution. 
This  is  due  to  the  fact  that  the  carbonate  combines  with 
the  carbonic  acid  present  in  the  water  to  form  calcium  acid 
carbonate,  which  is  soluble : 

CaCO,  +  H2C03  ^=>  Ca  (HCO3).2 

The  resulting  acid  carbonate  exists  only  in  solution,  since 
it  is  unstable  and  decomposes  into  the  normal  carbonate  on 
heating  or  on  evaporation  of  its  solution. 

Formation  of  caves.  Natural  waters  always  contain  more  or 
less  carbon  dioxide  in  solution.  In  the  case  of  certain  under- 
ground waters  the  amount  of  carbon  dioxide  is  comparatively 


FIG.  147.   A  crystal  of  Iceland  spar 


300  FIRST  COURSE  IN  CHEMISTRY 

large,  being  held  in  solution  by  pressure.  Such  waters  have 
a  marked  solvent  action  upon  limestone,  dissolving  both  the 
calcium  carbonate  and  the  magnesium  carbonate.  In  certain 
localities  this  solvent  action,  continued  through  geological  ages, 
has  resulted  in  the  formation  in  limestone  rock  of  large  caves, 
such  as  the  Mammoth  Cave  in  Kentucky.  Water  trickling 
through  the  roofs  of  these  caves  evaporates,  leaving  a  deposit 
of  calcium  carbonate  which,  as  the  process  continues,  often 

forms  icicle-shaped  masses, 
known  as  stalactites ;  or  the 
water  may  drip  upon  the  floor 
of  the  cave,  forming  similar 
masses  known  as  stalagmites 
(Fig.  148). 

Commercial  methods  for 
softening  water.  Hard  wa- 
ters contain  not  only  the  acid 
carbonates  of  magnesium  and 
calcium  but  also  the  chlo- 
rides and  sulfates  of  the 
metals,  together  with  small 

Quantities  of  other  salts.   On 
FIG.  148.    Stalactites  and 

stalagmites  a  small  scale  such  waters  are 

often    softened    by    adding 

borax  or  normal  sodium  phosphate.  These  precipitate  the 
calcium  and  magnesium  present,  at  the  same  time  leaving 
the  water  slightly  alkaline  and  therefore  adapted  for  clean- 
ing purposes.  On  a  large  scale,  as  in  the  case  of  a  city  water- 
supply,  the  water  is  softened  by  the  addition  of  calcium 
hydroxide  and  sodium  carbonate.  The  calcium  hydroxide 
converts  the  acid  carbonates  of  calcium  and  magnesium  into 
the  normal  carbonates,  which,  being  insoluble,  settle  out : 

CaHC03  +  Ca  (OH). >-  CaCO3  +  2  H2O 


THE  CALCIUM  FAMILY  301 

The  sodium  carbonate  used  converts  the  chlorides  and 
sulfates  of  calcium  and  magnesium  into  the  insoluble 
carbonates : 

CaS04  +  Na2C03  — >»  CaCO3  +  Na2SO4 

Water  softened  in  this  way  contains  sodium  sulfate  and 
chloride,  but  these  salts  are  not  objectionable. 

The  amounts  of  calcium  hydroxide  and  sodium  carbonate 
required  to  soften  any  given  water  are  calculated  from  an 
analysis  of  the  water.  These  are  added  to  the  water  in  large 
tanks  and  thoroughly  mixed.  The  mixture  is  then  run  out  into 
settling  basins.  After  the  solids  have  settled,  the  clear  water 
is  drawn  off. 

Calcium  oxide  (lime)  (CaO).  Pure  calcium  oxide  can  be 
prepared  by  burning  calcium  or  by  heating  its  nitrate  or 
carbonate.  The  more  or  less  impure  oxide  (commercially 
known  as  lime  or  quicklime)  is  prepared  by  strongly  heating 
limestone  in  large  furnaces  called  limekilns : 

CaCO, ^CaO  +  CO9 

o  -5 

When  pure,  lime  is  a  white  amorphous  substance. 
Heated  intensely,  as  in  the  oxy hydrogen  flame,  it  gives  a 
brilliant  light  called  the  limelight.  It  is  fusible  only  in  the 
heat  of  the  electric  furnace.  Water  acts  upon  lime  with 
the  evolution  of  a  great  deal  of  heat,  —  whence  the  name 
quicklime,  or  live  lime,  —  the  process  being  called  slaking. 
The  equation  is 

CaO  +  H2O *  Ca(OH)2  +  15,540  cal. 

Because  of  its  affinity  for  water  it  is  used  to  dry  gases.  It 
also  absorbs  carbon  dioxide,  forming  the  carbonate : 

CaO+C02 


302 


FIKST  COUKSE  IN  CHEMISTEY 


Lime  exposed  to  air  is  therefore  gradually  converted  into  the 
hydroxide  and  the  carbonate,  and  will  110  longer  slake  with 
water.  It  is  then  said  to  be  air-slaked.  Lime  is  produced  in 
enormous  quantities  for  use  in  making  calcium  hydroxide. 

Commercial  production  of  lime.  A  vertical  section  of  the  newer 
form  of  limekiln  is  shown  in  Fig.  149.  The  kiln  is  about  50  ft. 
in  height.  A  number  of  fire  boxes,  or  furnaces  A,  J,  are  built 
around  the  lower  part,  all  leading  into 
the  central  stack.  The  kiln  is  filled  with 
limestone  through  a  swinging  door  B. 
The  hot  products  of  combustion  are 
drawn  up  through  the  kiln,  and  the 
limestone  is  gradually  decomposed  by 
the  heat.  The  bottom  of  the  furnace 
is  so  constructed  that  a  current  of  air 
is  drawn  in  at  C,  and  this  serves  the 
double  purpose  of  cooling  the  hot  lime 
at  the  base  of  the  furnace  and  of  fur- 
nishing heated  oxygen  for  the  combus- 
tion. The  lime  is  dropped  into  cars  run 
under  the  furnace.  Generally  a  number 
of  these  kilns  are  operated  together,  as 
shown  in  Fig.  150. 
FIG.  149.  Cross  section 

of  a  modern  limekiln          Calcium    hydroxide    (slaked    lime) 
(Ca(OH)2).  This  compound  is  prepared 

by  adding  water  to  lime,  as  explained  above.  When  pure, 
it  is  a  light  white  powder.  It  is  sparingly  soluble  in  water, 
forming  a  solution  called  limewater.  Owing  to  its  cheapness 
it  is  used  in  the  industries  whenever  an  alkali  is  desired. 
It  is  used  in  the  preparation  of  ammonia,  bleaching  powder, 
and  the  hydroxides  of  sodium  and  potassium.  It  is  also  used 
to  remove  sulfur  compounds  and  carbon  dioxide  from  coal 
gas,  to  remove  the  hair  from  hides  in  making  leather,  for 
making  mortar  and  plaster,  and  for  liming  soils  (p.  110). 


THE  CALCIUM  FAMILY 


303 


Mortar  and  plaster.  Mortar  is  a  mixture  of  calcium  hydroxide 
and  sand.  When  it  is  exposed  to  the  air  or  spread  upon  porous 
materials,  moisture  is  removed  from  it  (partly  by  absorption  into 
the  porous  materials  and  partly  by  evaporation)  and  the  mortar 
becomes  firm,  or  sets.  At  the  same  time  carbon  dioxide  is  slowly 
absorbed  from  the  air  and  hard  calcium  carbonate  is  formed. 


Ca(OH)2  +  C02 


CaC0 


H20 


By  this  combined  action  the  mortar  becomes  very  hard  and 
adheres  firmly  to  the  surface  upon  which  it  is  spread.    The 

sand     serves     to     ___^ 

give  body  to  the 
mortar  and  makes 
it  porous.  It  also 
prevents  too  much 
shrinkage.  Plas- 
ter is  a  mixture 
of  calcium  hy- 
droxide and  hair, 
the  latter  serving 
to  hold  the  mass 
together. 

Bleaching  pow- 
der. When  chlo- 
rine   is    passed     FK,  15Q<  A  groui)  of  lilliekihls  in  a  modem  plant 
over  calcium  hy- 
droxide, there  is  formed  a  white  solid  compound  having 
the  formula  CaOCl2,  and  known  as  bleaching  powder,  chloride 
of  lime,  or  simply  bleach : 

Ca(OH)2  +  C12 >-  CaOCl2  4-  H2O 

When  this  compound  is  treated  with  an  acid,  chlorine  is 
evolved : 

I-  H  SO. >-  CaSO.  +  HO  4-  01 


304 


FIRST  COURSE  IN  CHEMISTRY 


When  exposed  to  the  air,  bleaching  powder  is  slowly 
acted  upon  by  moisture  and  carbon  dioxide,  with  the 
liberation  of  hypochlorous  acid  (HC1O),  which  is  a  good 
disinfectant. 

Bleaching  powder  is  prepared  in  large  quantities  for  use 
as  a  bleaching  agent,  as  a  disinfectant,  and  as  an  agent  for 
purifying  city  water-supplies. 

In  the  commercial  preparation  of  bleaching  powder  the 
calcium  hydroxide  is  spread  to  a  depth  of  2  or  3  inches  upon 
the  floor  of  a  room,  usually  made  of  lead  or  concrete  (Fig.  151). 
The  chlorine  enters  near  the  top  at  A.  Any  unabsorbed  chlo- 
rine passes  out  at  B  and  into  the  adjoining  chamber  at  C. 


g-*"  ~"E 

FIG.  1  jl.    Chambers  for  the  manufacture  of  bleaching  powder 

Calcium  sulfate  (CaSOJ.  Calcium  surf  ate  occurs  in  nature 
in  several  different  forms,  the  most  common  of  which  is 
gypsum  (CaSO4'2  H2O).  This  is  quarried  in  large  amounts 
in  New  York,  Michigan,  and  Oklahoma.  It  is  used  as  a  filler 
in  making  paper  (p.  227),  as  a  constituent  of  fertilizers, 
and  especially  in  making  plaster  of  Paris.  It  is  but  slightly 
soluble  in  water. 

Plaster  of  Paris  ((CaS04)2  •  H20).  This  is  a  fine  white 
powder  obtained  by  carefully  heating  gypsum.  When 
water  is  added  to  the  powder,  a  plastic  mass  is  formed  which 
quickly  hardens,  or  sets.  This  property  makes  it  valuable 
for  molding  casts,  for  stucco  work,  and  for  a  finishing  coat 
on  plastered  walls.  Broken  bones  are  often  held  in  place  by 
casts  of  plaster  of  Paris  until  they  grow  together. 


THE  CALCIUM  FAMILY 


305 


Calcium  carbide  (CaC2).  This  compound  is  prepared  on  a 
large  scale  for  use  in  the  manufacture  of  acetylene  (p.  207) 
and  in  making  fertilizers.  It  is  made  by  heating  a  mixture 
of  lime  and  coke  in  an  electric  furnace : 


CaO  +  3  C 


CaO  +  CO 


The  pure  carbide  is  a  colorless,  transparent  solid.  The 
commercial  article  is  a  dull-gray  porous  substance  which 
contains  many  impurities.  It 
is  placed  on  the  market  in 
air-tight  containers. 

Commercial  preparation.  While 
calcium  carbide  was  first  pre- 
pared in  1836,  it  was  not  until 
1893  that  it  became  a  commer- 
cial product.  The  general  princi- 
ples involved  in  its  preparation 
are  illustrated  in  Fig.  152,  which 
represents  a  simple  type  of  a 
carbide  furnace.  The  base  of 
the  furnace  is  provided  with  a 
large  block  of  carbon  A ,  which 
serves  as  one  of  the  electrodes. 
The  other  electrodes,  B,  B,  sev- 
eral in  number,  are  arranged 

horizontally  at  some  distance  above  this.  A  mixture  of  coal  and 
lime  is  fed  into  the  furnace  through  the  trap  top  C,  and  in  the 
lower  part  of  the  furnace  this  mixture  becomes  intensely  heated, 
forming  liquid  carbide.  This  is  drawn  off  through  the  tap  hole  D. 

Calcium  cyanamide  (CaN2C).  When  nitrogen  is  passed 
over  hot  calcium  carbide,  the  two  react,  forming  a  com- 
pound known  as  calcium  cyanamide : 


FIG.  152.   A  furnace  for  the  manu- 
facture of  calcium  carbide 


CaN  C  +  C 


306  FIRST  COURSE  IN  CHEMISTRY 

The  commercial  product  is  impure,  containing  carbon,  lime, 
besides  about  60  per  cent  of  the  cyanamide.  This  product 
is  known  as  lime-nitrogen.  This  is  ground,  mixed  with  water 
(which  slakes  the  lime),  and  in  this  form  sold  as  a  fertilizer 
under  the  name  cyanamide.  Its  value  as  a  fertilizer  lies  in 
the  fact  that  all  of  its  nitrogen  is  available  as  a  plant  food. 

The  nitrogen  used  in  the  preparation  of  cyanamide  is  ob- 
tained by  passing  air  (freed  from  moisture  and  carbon  dioxide) 
over  hot  copper  (p.  81).  By  means  of  this  compound,  there- 
fore, it  becomes  possible  to  utilize  the  nitrogen  in  the  air  as  a 
plant  food  (p.  83). 

Calcium  phosphate  (Ca3(P04)2).  Tins  important  substance 
occurs  in  nature  as  the  mineral  phosphorite  and  as  a  con- 
stituent of  apatite.  Large  amounts  of  it  occur  in  the  form 
of  rock  phosphate,  which  is  found  especially  in  Florida, 
Tennessee,  and  some  of  the  Western  states.  It  is  also  the 
chief  mineral  constituent  of  bones.  Bone  ash  is,  therefore, 
nearly  pure  calcium  phosphate. 

Other  compounds  of  calcium.  Calcium  chloride  (CaCl2)  occurs 
in  sea  water  and  is  formed  in  large  quantities  as  a  by-product 
in  the  Solvay  process  for  making  sodium  carbonate.  The  an- 
hydrous salt  readily  absorbs  moisture  and  is  used  as  an  agent 
for  drying  gases.  A  solution  of  the  salt  is  used  as  a  brine  in 
the  manufacture  of  ice  (p.  124).  It  has  also  been  used  to  lay 
the  dust  on  roads,  and  mines  have  been  sprinkled  with  it 
in  the  hope  of  preventing  dust  explosions.  Calcium,  fluoride 
(CaF2)  occurs  in  nature  in  the  form  of  fluorlte.  It  is  mined  in 
large  quantities,  especially  in  Illinois,  and  is  used  in  the  prep- 
aration of  hydrofluoric  acid,  in  the  manufacture  of  opaque 
glass,  and  in  various  metallurgical  operations.  Calcium  sul- 
fide  (CaS)  is  a  by-product  in  the  Leblanc  process  for  making 
sodium  carbonate.  The  commercial  salt  is  sometimes  used  as 
a  luminous  paint,  since,  after  exposure  to  a  bright  light,  it 
will  glow  in  the  dark.  Calcium  acid  s-ulfite  (Ca(HS03)2)  is 


THE  CALCIUM  FAMILY  SOT 

used  as  a  preservative,  and  in  large  quantities  in  the  manufac- 
ture of  paper  (p.  227).  A  number  of  calcium  silicates  are 
known,  and  derive  their  chief  interest  from  the  fact  that  they 
are  important  constituents  of  cement  and  glass. 


STRONTIUM  AND  BARIUM 

General.  These  elements  are  much  rarer  than  calcium, 
are  difficult  to  prepare,  and  have  no  commercial  uses.  Their 
compounds  resemble  those  of  calcium  in  composition  and 
properties.  Strontium  compounds,  especially  the  nitrate, 
when  ignited  with  oxidizable  substances,  give  a  brilliant 
crimson  color,  and  on  this  account  are  used  in  the  manu- 
facture of  red  lights.  Under  similar  conditions  barium  ni- 
trate gives  a  green  light.  The  following  compounds  of 
barium  are  of  importance. 

Oxides  of  barium.  Barium  oxide  (BaO)  can  be  obtained 
by  strongly  heating  the  nitrate : 

2  Ba(NO3)2 >-  2  BaO  4-  4  NO2  +  O2 

Heated  to  a  low  red  heat  in  the  air,  the  oxide  combines 
with  oxygen,  forming  the  peroxide  (BaO0),  which  is  used 
in  making  hydrogen  peroxide  (p.  52). 

Barium  chloride  (BaCl2  •  2  H20).  Barium  chloride  is  a 
white  solid.  It  is  used  in  the  laboratory  as  a  reagent  to 
detect  the  presence  of  sulfuric  acid  or  soluble  sulfates, 
reacting  with  these  to  form  the  insoluble  barium  sulfate. 

Barium  sulfate  (barite)  (BaSOJ.  Barium  sulfate  occurs 
in  nature  as  a  heavy  white  mineral  known  as  barite.  It  is 
precipitated  as  a  crystalline  powder  when  a  barium  salt  is 
added  to  a  solution  of  a  sulfate  or  to  sulfuric  acid : 

BaCl2  +  H2S04  — +  BaS04  +  2  HC1 
It  is  used  in  large  quantities  in  the  manufacture  of  paints. 


308  FIEST  COURSE  IN  CHEMISTRY 

EXERCISES 

1.  What  properties  have  the  alkaline  earth  metals  in  common 
with  the  alkali  metals?    In  what  respect  do  they  differ? 

2.  Write  the  equation  for  the  reaction  between  calcium  carbide 
and  water. 

3.  How  is  calcium  chloride  removed  from  hard  water? 

4.  Would  air-slaked   lime  do   for  making  mortar?    Would  it 
serve  for  liming  acid  soils  (p.  110)? 

5.  Why  would  you  expect  calcium  carbide  to  contain  impurities  ? 

6.  How  do  you  explain  the  fact  that  calcium  carbonate  can  be 
decomposed  into  calcium  oxide  and  carbon  dioxide,  and  yet  calcium 
oxide  absorbs  carbon  dioxide  from  the  air  to  form  the  carbonate  ? 

7.  Could  barium  hydroxide  be  used  in  place  of  calcium  hydrox- 
ide in  testing  for  carbon  dioxide  ? 

8.  Calcite  and  gypsum  often  resemble  each  other  in  appearance. 
How  could  you  easily  distinguish  between  the  two  ? 

9.  What  weight  of  plaster  of  Paris  can  be  made  by  heating 
1  ton  of  gypsum?   Ans.  1686  Ib. 

10.  What  weight  of  limestone  is  necessary  to  prepare  10  tons 
of  lime?  Ans.  39,262  Ib. 

11.  What  weight  of  water  is  necessary  to  slake  1  ton  of  lime? 
Ans.  642  Ib. 

12.  How  could  you  prove  that  dried  mortar  contains  calcium 
carbonate  and  sand? 

TOPICS  FOR  THEMES 

If  possible,  inspect  a  limekiln  and  write  a  description  of  the 
process. 

The  great  caves  of  the  United  States  (see  encyclopedia). 

Find  out  whether  any  of  the  factories  in  your  vicinity  soften  the 
water  used  in  their  boilers,  and  if  they  do,  write  a  description  of 
the  process. 


CHAPTER  XXXII 
FERTILIZERS 

Plant  food;  fertilizers.  With  the  exception  of  carbon 
dioxide  (and  possibly  a  little  oxygen)  absorbed  from  the 
air,  the  growing  plant  derives  its  nourishment  from  the 
soil.  In  order  that  vegetation  may  thrive,  it  is  essential, 
therefore,  that  the  soil  should  contain  an  adequate  supply 
of  appropriate  plant  food.  Moreover,  since  this  supply  is 
continually  being  drawn  upon  by  the  growing  plant,  it  is 
necessary,  in  order  that  the  soil  may  retain  its  fertility, 
that  the  ingredients  so  withdrawn  shall  be  returned  to  it. 
It  is  for  this  purpose  that  fertilizers  are  used. 

Constituents  of  fertilizers.  While  a  number  of  elements  are 
essential  to  the  growth  of  the  plant,  experience  has  shown 
that  in  general  the  fertility  of  the  soil  may  be  maintained  by 
adding  three  substances :  (1)  nitrogenous  matter,  (2)  phos- 
phates of  calcium,  and  (3)  compounds  of  potassium. 

Sources  of  fertilizers.  The  commercial  sources  of  each 
of  the  constituents  of  fertilizers  are  as  follows: 

1.  Nitrogenous  matter.    This  is  obtained  from  a  number 
of  sources :  sodium  nitrate,  ammonium  sulfate,  and  cyana- 
mide  ;  also  nitrogenous  organic  matter,  such  as  dried  blood, 
the  waste  from  slaughterhouses,   and,   especially,   animal 
excrements. 

2.  Phosphates.     Ground  bones  are   especially  valuable, 
since  they  contain  some  nitrogen  in  addition  to  calcium 
phosphate.    This  source,  however,  is  entirely  inadequate, 
and  the  great  supply  comes  from  the  rock  phosphates, 

309 


310  FIRST  COURSE  IN  CHEMISTRY 

which  contain  about  70  per  cent  calcium  phosphate.  These 
rock  phosphates  are  quarried  in  large  quantities,  especially  in 
Florida  (Fig.  153)  and  Tennessee.  Since  calcium  phosphate 
is  nearly  insoluble,  the  rock  is  ground  and  then  treated  with 
sulfuric  acid.  This  converts  the  insoluble  calcium  phosphate 
into  the  soluble  calcium  acid  phosphate  (CaH4(PO4)2) : 

Caa(P04)2  +  2  H2S04  — >-  2  CaSO4  +  CaH4(PO4)2 

The  calcium  sulfate  also  adds  to  the  value  of  the  fertilizer, 
furnishing  sulfur  and  improving  the  physical  qualities  of 


the  soil.  Certain  products  (slags)  formed  in  the  manu- 
facture of  steel  contain  phosphorus,  and  are  also  used  in 
fertilizers. 

3.  Potassium  compounds.  These  are  obtained  from  the 
Stassfurt  mines.  Kainite  (KC1  •  MgSO4  .  3  H2O)  is  the 
most  common  of  the  minerals  used  (p.  286).  Wood  ashes 
are  excellent,  but  the  supply  is  limited. 

Commercial  fertilizers.  As  a  rule  the  fertilizers  on  the 
market  are  mixtures  of  the  three  fundamental  materials 
referred  to  above.  The  composition  is  varied  according  to 


FERTILIZERS 


311 


the  crop  to  be  grown  as  well  as  to  the  nature  of  the  soil ; 
for  example,  potatoes  demand  a  fertilizer  rich  in  potassium, 
while  a  cereal,  such  as  wheat,  is  benefited  more  by  one 
rich  in  phosphates.  Instead  of  using  a  fertilizer  containing 
all  three  constituents,  many  farmers  prefer  to  find  out  by 
experiment  just  what  plant  food  is  lacking  in  their  soils 
and  then  to  make  a  proper  mixture  of  such  fertilizing 
materials  as  will  furnish  the  desired  food.  Fig.  154  shows 
the  result  of  such  an  ex- 
periment. Pot  1  shows  the 
result  obtained  with  no  fer- 
tilizer ;  pot  2,  with  one  com- 
bination ;  and  pot  3,  with 
another. 

The  liming  of  soils.  Some- 
times a  soil  may  be  unproduc- 
tive from  other  causes  than 
deficiency  in  plant  food.  For 
example,  it  may  be  acid  in 
reaction,  and  in  such  a  case  it 
may  be  made  fertile  once  more 
by  being  timed  (p.  110). 


FIG.  154.    The  effect  of  fertilizers 
upon  the  growth  of  plants 


The  utilization  of  atmospheric  nitrogen.  It  has  been 
pointed  out  that  with  few  exceptions  plants  have  not  the 
power  of  assimilating  free  nitrogen  (p.  83).  Moreover,  it 
is  inevitable  that  the  supplies  of  sodium  nitrate  and  ammo- 
nium sulfate,  which  are  now  the  chief  nitrogenous  prod- 
ucts used  in  the  manufacture  of  commercial  fertilizers,  will 
sooner  or  later  become  exhausted.  Repeated  efforts  have 
therefore  been  made  to  utilize  the  inexhaustible  supply  of 
nitrogen  in  the  atmosphere.  It  has  been  found  possible  to 
do  this  by  converting  the  nitrogen  into  compounds  which 
contain  the  element  in  a  form  available  for  plant  food. 


312  FIRST  COURSE  IN  CHEMISTEY 

The  following  methods  may  be  used  for  effecting  this 
change  :  (1)  the  nitrogen  may  be  converted  into  cyanamide 
(p.  305)  ;  (2)  the  nitrogen  may  be  converted  into  nitric 
acid  by  electric  sparking,  and  then  into  nitrates  (p.  128)  ; 
(3)  ammonia  may  be  formed  by  heating  a  mixture  of  nitro- 
gen and  hydrogen  (p.  123),  and  this,  with  sulfuric  acid, 
gives  ammonium  sulfate.  These  methods  are  now  used 
commercially  to  a  limited  extent,  and  it  appears  certain 
that  before  long  the  quantity  of  nitrogenous  compounds 
annually  required  for  fertilizers  jwill  be  prepared  from 
atmospheric  nitrogen. 

EXERCISES 

1.  If  the  crops  grown  on  a  soil  were  not  removed,  would  the  soil 
diminish  in  fertility  ?    Would  it  increase  in  fertility  ? 

2.  Why  does  treatment  of  bones  with  sulfuric  acid  make  them 
more  available  as  plant  food? 

3.  In  the  preparation  of  fertilizers,  what  weight  of  sulfuric  acid 
containing  50  per  cent  hydrogen  sulfate  is  necessary  in  the  treatment 
of  10  tons  of  rock  phosphate  containing  70  per  cent  pure  calcium 
phosphate?    Ans.  17,702  Ib. 

TOPIC  FOR  THEMES 

The  utilization  (fixation)  of  atmospheric  nitrogen  (Duncan, 
Chemistry  of  Commerce). 


CHAPTER  XXXIII 
THE  MAGNESIUM  FAMILY 


NAME 

SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

MKLTING 
POINT 

BOILING 
POINT 

OXIDE 

Magnesium 
Zinc  . 

Mg 

Zn 

24.32 
65.37 

1.74 

7.10 

651° 
419.4° 

920° 
950° 

MgO 

ZnO 

Cadmium   . 

Cd 

112.4 

8.64 

320.9° 

778° 

CdO 

The  family.  In  the  magnesium  family  are  included  the 
four  elements:  magnesium,  zinc,  cadmium,  and  mercury. 
Between  the  first  three  of  these  metals  there  is  a  close 
family  resemblance.  Mercury  in  some  respects  is  more 
similar  to  copper  and  will  be  studied  in  connection  with 
that  metal. 

MAGNESIUM 

Occurrence.  Magnesium  is  a  very  abundant  element  in 
nature,  ranking  a  little  below  calcium  in  this  respect.  Like 
calcium,  it  is  a  constituent  of  many  rocks  and  also  occurs 
in  the  form  of  soluble  salts.  It  is  a  constituent  of  chlorophyll 
and  is  therefore  essential  to  plant  life. 

Preparation.  The  metal  magnesium,  like  most  metals 
whose  oxides  are  difficult  to  reduce  with  carbon,  is  made 
by  electrolysis ;  but  instead  of  using  as  the  electrolyte  the 
melted  anhydrous  chloride,  which  is  difficult  to  obtain,  the 
mineral  carnallite  (p.  286)  is  used.  This  is  melted  in  an 
iron  pot  which  also  serves  as  the  cathode  in  the  electrolysis. 

313 


314  FIRST  COUKSE  IN  CHEMISTRY 

A  rod  of  carbon  dipping  into  the  melted  salt  serves  as  the 
anode.  The  apparatus  is  very  similar  to  those  employed  in 
the  preparation  of  sodium  and  calcium. 

Properties.  Magnesium  is  a  silvery-white  metal  of  small 
density.  It  is  usually  sold  in  the  form  of  thin  ribbon  or  of 
wire  or  as  a  powder.  Air  does  not  act  rapidly  upon  it,  but 
a  thin  film  of  basic  carbonate  forms  upon  its  surface,  dim- 
ming its  bright  luster.  The  common  acids  dissolve  it  with 
the  formation  of  the  corresponding  salts.  It  can  be  ignited 
readily,  and  in  burning  it  gives  a  brilliant  white  light. 
This  light  is  very  rich  in  the  rays  which  affect  photo- 
graphic plates,  and  the  metal,  in  the  form  of  tine  powder, 
is  extensively  used  in  the  production  of  flash  lights  and  for 
white  lights  in  pyrotechnic  displays.  When  used  for  this 
purpose,  the  powder  is  mixed  with  an  oxidizing  agent, 
potassium  chlorate  being  the  one  commonly  employed. 

Magnesium  oxide  (magnesia)  (MgO).  Magnesium  oxide, 
sometimes  called  magnesia  or  magnesia  usta,  resembles  lime 
in  many  respects.  It  is  much  more  easily  formed  than  lime 
and  can  be  made  in  the  same  way,  namely,  by  heating  the 
carbonate.  It  is  a  white  powder,  very  soft  and  bulky,  and 
is  unchanged  by  heat  even  at  very  high  temperatures.  For 
this  reason  it  is  used  in  the  manufacture  of  crucibles,  for 
lining  furnaces,  and  for  other  purposes  where  a  refractory 
basic  substance  is  needed. 

Magnesium  hydroxide  (Mg(OH)2).  The  hydroxide  of 
magnesium  is  but  slightly  soluble  in  water  and  can  be  pre- 
cipitated by  adding  a  soluble  base  to  a  salt  of  magnesium : 

MgCl2  +  Ca(OH)2 >  Mg(OH)2  +  CaCl, 

It  dissolves  sufficiently  to  give  a  slightly  alkaline  reaction, 
and  is  a  moderately  strong  base.  It  is  a  white  amorphous 
substance,  and  is  converted  into  the  oxide  when  heated. 


THE  MAGNESIUM  FAMILY  315 

Magnesium  carbonate  (MgC03).  Magnesium  carbonate 
occurs  in  a  number  of  localities  as  magnetite,  which  is 
usually  amorphous,  but  sometimes  forms  pure  crystals 
resembling  calcite.  More  frequently  it  is  found  associated 
with  calcium  carbonate.  The  mineral  dolomite  has  the  com- 
position CaCO3  •  MgCO3.  Limestone  containing  smaller 
amounts  of  magnesium  carbonate  is  known  as  dolomitic 
limestone.  Dolomite  is  one  of  the  most  common  rocks, 
forming  whole  mountain  masses.  It  is  harder  and  less 
readily  attacked  by  acids  than  limestone.  It  is  valuable 
as  a  building-stone  and  for  foundations  and  as  ballast  for 
roadbeds.  Like  calcium  carbonate,  magnesium  carbonate  is 
insoluble  in  water,  but  readily  dissolves  in  water  containing 
carbon  dioxide,  forming  the  acid  carbonate : 

MgCO,  +  H,0  +  C02  — >  Mg(HCO.), 

Boiler  scale.  When  water  which  contains  certain  salts  in 
solution  is  evaporated  in  steam  boilers,  a  hard,  insoluble  mate- 
rial called  scale  deposits  in  the  boiler.  The  formation  of  this 
scale  may  be  due  to  several  distinct  causes : 

1.  To  the  deposit  of  calcium  sulfate.    This  salt,  while  sparingly 
soluble  in  cold  water,  is  almost  completely  insoluble  in  super- 
heated   water.     Consequently   it    is   precipitated  when   water 
containing  it  is  heated  in  a  boiler. 

2.  To  decomposition  of  acid  carbonates.    As  we  have  seen,  cal- 
cium and  magnesium  acid  carbonates  are  decomposed  on  heat- 
ing, forming  insoluble  normal  carbonates  : 

Ca(HC03)2 *  CaC03  +  HaO  +  C02 

3.  To  hydrolysis  of  magnesium  salts.    Magnesium  chloride,  and 
to  some  extent  magnesium  sulfate,  undergo  hydrolysis  when 
superheated  in  solution,  and  the  magnesium  hydroxide,  being 
sparingly  soluble,  precipitates : 

MgCl2  +  2  HaO *  Mg(OH)2  +  2  HC1 


316 


FIRST  COURSE 


CHEMISTRY 


This  scale  adheres  tightly  to  the  boiler  tubes  in  compact  layers 
(Fig.  155),  and,  being  a  nonconductor  of  heat,  causes  much 
waste  of  fuel.  It  is  very  difficult  to  remove,  owing  to  its  hard- 
ness and  its  resistance  to  reagents.  Thick  scale  sometimes 
cracks,  and  the  water  coming  in  contact  with  the  overheated 
iron  occasions  an  explosion. 

Magnesium  sulfate  (MgSOJ.  Like  the  chloride,  magne- 
sium sulfate  is  found  rather  abundantly  in  springs  and 
in  salt  deposits.  Deposits  of  the  almost  pure  solid  salt 

having  the  composition 
MgSO4.7H2Ohavebeen 
found  in  Wyoming  and 
Washington.  It  is  often 
called  Epsom  salt  be- 
cause of  its  occurrence  in 
the  waters  of  the  Epsom 
springs  in  England. 

Magnesium  sulfate  is 
used  to  a  small  extent 
in  the  preparation  of 
sodium  and  potassium 
sulfates;  for  weighting 
cotton  cloth  in  the  dye 
industry  ;  in  tanning  ; 
and  in  the  manufacture  of  paints  and  laundry  soaps.  To 
some  extent  it  is  used  in  medicine. 

Magnesium  silicates.  Many  silicates  containing  magne- 
sium are  known,  and  some  of  them  are  important  substances. 
Serpentine,  asbestos,  talc  (or  soapstone),  and  meerschaum  are 
examples  of  such  substances.  Asbestos  is  soft  and  fibrous 
and  a  nonconductor  of  heat.  It  is  used  for  fireproof  material 
in  a  great  variety  of  forms,  such  as  cloth,  paper,  board,  and 
rope.  It  is  also  used  as  a  covering  for  pipes,  furnaces,  and 


FIG.  155.    Cross  section  of  a  boiler  tube 
showing  the  deposit  of  boiler  scale 


THE  MAGNESIUM  FAMILY  317 

boilers,  to  diminish  heat  radiation.  Soapstone  is  valuable 
for  sinks  and  table  tops,  and  in  finely  ground  form  as  a 
toilet  powder  and  foot  ease,  and  is  called  French  chalk. 
Meerschaum  is  used  for  pipe  bowls  and  similar  articles. 

ZINC 

Occurrence.  Zinc  never  occurs  free  in  nature.  It  is  not 
a  constituent  of  common  rocks  and  minerals,  and  its  occur- 
rence is  rather  local  and  confined  to  definite  deposits  or  to 
pockets.  It  occurs  chiefly  in  the  following  ores :  sphalerite 
(zinc  blende)  (ZnS);  zincite  (ZnO);  smithsonite  (ZnCO3); 
franklinite  (ZnO  •  Fe0Og).  One  fourth  of  the  world's  out- 
put of  zinc  comes  from  the  United  States  —  Missouri,  Kan- 
sas, and  New  Jersey  being  the  largest  producers. 

Metallurgy.  The  ores  employed  in  the  preparation  of 
zinc  are  chiefly  the  sulfide,  oxide,  and  carbonate.  They  are 
first  heated  in  the  air,  by  which  process  they  are  changed 
into  the  oxide,  this  process  being  called  roasting : 

ZnCO3 >-  ZnO  +  CO0 

2  ZnS  +  3  O9 >-  2  ZnO  +  2  SO 

I  2 

The  oxide  is  then  mixed  with  coal  dust,  and  the  mixture 
is  heated  in  earthenware  retorts.  The  oxide  is  reduced  by 
this  means  to  the  metallic  state,  and  the  zinc,  being  heated 
above  its  boiling  point,  distills  and  is  collected  in  suitable 
receivers  and  is  drawn  off  into  molds.  In  this  form  it  is 
called  spelter.  Commercial  zinc  often  contains  a  number  of 
impurities,  especially  carbon,  arsenic,  and  iron. 

Properties.  Pure  zinc  is  a  rather  heavy  bluish-white 
metal  with  a  high  luster.  It  melts  at  about  420°,  and  if 
heated  much  above  this  temperature  in  the  air,  it  takes  fire 
and  burns  with  a  bluish  flame.  It  boils  at  about  950°. 


318  FIRST  COURSE  IN  CHEMISTRY 

Many  of  the  properties  of  zinc  are  much  influenced 
by  the  temperature  and  previous  treatment  of  the  metal. 
When  cast  into  ingots  from  the  liquid  state,  it  becomes  at 
ordinary  temperatures  quite  hard,  brittle,  and  highly  crys- 
talline. At  100°-150°  it  is  malleable  and  can  be  rolled 
into  thin  sheets ;  at  higher  temperatures  it  again  becomes 
very  brittle.  When  once  rolled  into  sheets  it  retains  its 
softness  and  malleability  at  ordinary  temperatures.  When 
melted  and  poured  into  water  it  forms  thin,  brittle  flakes, 
and  in  this  condition  is  called  granulated  zinc  or  mossy  zinc. 

Zinc  is  tarnished  superficially  by  moist  air,  but  beyond 
this  is  not  affected  by  it.  When  the  metal  is  quite  pure, 
sulfuric  and  hydrochloric  acids  act  upon  it  very  slowly ; 
when,  however,  it  contains  small  amounts  of  other  metals, 
such  as  magnesium  or  copper,  or  when  it  is  merely  in  con- 
tact with  another  metal,  brisk  action  takes  place  and  hy- 
drogen is  evolved.  For  this  reason,  when  pure  zinc  is  used  in 
the  preparation  of  hydrogen,  a  few  drops  of  copper  sulfate 
are  often  added  to  the  solution  to  assist  the  chemical  action. 

Uses  of  zinc.  The  chief  use  of  zinc  is  in  the  manufacture 
of  galvanized  iron.  This  is  sheet  iron  or  wire  covered  with 
a  thin  layer  of  zinc,  which  protects  the  iron  from  rusting. 
About  two  thirds  of  all  the  zinc  produced  is  used  in  this 
way.  Sheet  zinc  is  used  as  a  lining  for  sinks  and  water- 
containers.  Large  quantities  of  the  metal  are  used  in  mak- 
ing brass  and  other  alloys  (p.  360),  in  the  construction  of 
electrical  batteries,  and  in-  separating  silver  from  lead 
(p.  373).  In  the  laboratory  it  is  used  in  the  preparation  of  hy- 
drogen and,  in  the  form  of  zinc  dust,  as  a  reducing  agent. 

Manufacture  of  galvanized  iron.  Fig.  156  shows  the  method 
used  in  making  galvanized  iron.  The  plates  of  iron  pass  under 
the  rollers  at  A  and  on  into  the  pot  of  melted  zinc  B.  The 
zinc  adheres  to  tlie  iron,  and  the  resulting  plate  is  passed  under 


THE  MAGNESIUM  FAMILY  319 

the  roller  C  to  remove  the  excess  of  zinc  and  to  render  the 
surface  smooth.  Sometimes  the  zinc  is  deposited  on  the  iron 
by  electrolytic  methods. 

Zinc  oxide  (zinc  white)  (ZnO).  Zinc  oxide  occurs  in  im- 
pure form  in  nature,  being  colored  red  by  manganese  and 
iron  compounds.  It  can  be  prepared  in  the  same  way  as 
magnesium  oxide,  namely,  by  heating  zinc  carbonate  or 
hydroxide,  but  is  more  often  made  by  burning  the  metal. 


FIG.  150.    The  manufacture  of  galvanized  sheet  iron 

Zinc  oxide  is  a  pure-white  powder  which  is  much  used 
as  a  white  pigment  in  paints,  under  the  name  of  zinc  white. 
It  has  an  advantage  over  white  lead  in  that  it  is  not 
changed  in  color  by  sulfur  compounds,  while  lead  com- 
pounds turn  black.  Many  thousand  tons  of  zinc  oxide  are 
used  in  paints  each  year.  It  is  also  used  as  a  filler  in  the 
manufacture  of  rubber  goods.  More  than  two  thousand  tons 
are  used  annually  in  the  manufacture  of  automobile  tires. 


320  FIRST  COURSE  IN  CHEMISTRY 

Zinc  sulfate  (ZnSOJ.  This  salt  is  readily  crystallized 
from  concentrated  solutions  in  transparent  colorless  crystals 
which  have  the  formula  ZnSO4  •  7  H2O  and  are  called  white 
vitriol.  It  is  prepared  commercially  by  careful  roasting  of 
thesulflde:  ZnS  +  2  O  —  ^  ZnSO 


Zinc  chloride  (ZnCl2).  This  salt  is  very  soluble  in  Avater 
and  has  a  strongly  acid  reaction.  It  has  germicidal  prop- 
erties and  is  used  to  preserve  railroad  ties  and  other 
wooden  timbers  especially  subject  to  decay. 

Zinc  sulfide  (ZnS).  This  substance  occurs  as  the  mineral 
sphalerite,  and  is  one  of  the  most  valued  ores  of  zinc. 
Very  large  deposits  occur  in  southwestern  Missouri.  The 
natural  mineral  is  found  in  large  crystals  or  masses,  resem- 
bling resin  in  color  and  luster.  It  is  insoluble  in  water  and, 
when  prepared  by  precipitation,  is  white.  Lithopone  is  a 
mixture  of  the  two  solids,  barium  sulfate  and  zinc  sulfide, 
made  by  bringing  together  barium  sulfide  and  zinc  sulfate  in 

solution:        T,  0      r/   ^^  rr  0 

BaS  4-  ZnSO4  -  >•  BaSO4  +  ZnS 

It  is  a  valuable  white-paint  pigment. 

Preservation  of  wood.  With  the  rapid  disappearance  of  the 
forests  the  preservation  of  wood  from  decay  (fungous  growths) 
becomes  a  very  important  problem.  When  the  wood  is  to  be 
exposed  merely  to  atmospheric  conditions,  it  is  preserved  by 
paints  and  varnishes.  When  it  must  be  partly  buried  in  the 
ground  (railway  ties,  fence  posts),  it  is  treated  with  germicidal 
preservatives.  Those  most  frequently  used  are  zinc  chloride, 
copper  sulfate,  and  creosote. 

The  wood  is  placed  in  closed  boilers  in  baths  of  the  appro- 
priate liquid,  and  the  air  is  exhausted  so  that  the  liquid  may 
be  more  readily  driven  into  the  pores  of  the  wood.  Frequently 
the  latter  process  is  assisted  by  the  application  of  considerable 
pressure  to  the  liquid  after  the  air  has  been  pumped  out. 


THE  MAGNESIUM  FAMILY  321 

Cadmium.  This  element  very  closely  resembles  zinc  in 
most  respects.  Some  of  its  alloys  are  characterized  by  having 
low  melting  points.  Its  compounds  are  similar  in  composi- 
tion to  the  corresponding  ones  of  magnesium  and  zinc. 

EXERCISES 

1.  What  metals  already  studied  are  prepared  by  electrolysis? 

2.  When  acids  act  upon  magnesium,  is  hydrogen  liberated? 

3.  What  property  of  magnesium  was  taken  advantage  of  in  the 
isolation  of  argon? 

4.  How    could    you    distinguish    between    Glauber's    salt    and 
Epsom  salt? 

5.  Account  for  the  fact  that  paints  made  of  y.inc  oxide  are  not 
colored  by  hydrogen  sulfide. 

6.  Why  does  not  zinc  occur  in  the  uncombined  state  in  nature? 

7.  What  reaction  takes  place  when  zinc  is  added  to  a  solution 
of  copper  sulfate?    (See  displacement  series.) 

8.  Write  equations  showing  how  the  following  compounds  of 
zinc  may  be  obtained  from  metallic  zinc  :  the  oxide,  chloride,  nitrate, 
carbonate,  sulfate,  sulfide,  hydroxide. 

9.  AVhat  is  the  composition  of  a  lime  made  from  dolomite? 
10.    Which  would  yield  the  most  zinc,  1  ton  of  sphalerite  or  1  ton 

of  franklinite? 

TOPICS  FOR  THEMES 

The  zinc  mines  of  Missouri  (United  States  Geological  Survey, 
AVashington). 

The  galvanizing  of  iron  (see  encyclopedia). 


CHAPTER  XXXIV 


ALUMINIUM 

Occurrence.    Aluminium  never  occurs  in  the  free  state  in 
nature,  owing  to  its  great  affinity  for  oxygen.   In  combined 

form,  as  oxide,  silicates,  and 
a  few  other  compounds,  it  is 
both  abundant  and  widely 
distributed,  being  an  essen- 
tial constituent  of  all  soils, 
and  of  most  rocks  except 
limestone  and  sandstone.  It 
is  estimated  that  aluminium 
constitutes  about  8  per  cent 
of  the  earth's  crust.  Cry- 
olite (Na3AlF6),  found  in 
Greenland,  and  bauxite, 
which  is  an  aluminium  hy- 
droxide usually  mixed  with 
some  iron  hydroxide,  are 
important  minerals.  In  the 
industries  the  metal  is  called 
aluminum,  but  its  chemical 
name  is  aluminium. 

Preparation.     Aluminium 
was  first  prepared  by  Wohler, 
in  1827,  by  heating  anhydrous  aluminium  chloride  with 
potassium  : 


Courtesy  of  The  Chemists'  Club,  Xew  York 

FH..  157.    Charles  Martin  Hall 

(1863-1914) 

The  American  chemist  who  developed 

the  electrical  method  for  producing 

aluminium 


A1CL  +  3  K 


3  KC1  +  Al 


322 


ALUMINIUM 


323 


Although  the  metal  is  very  abundant  in  nature  and  pos- 
sesses many  desirable  properties,  the  cost  of  separating  it 
from  its  ores  was  so  great  that  it  remained  almost  a  curi- 
osity until  comparatively  recent  years.  With  the  develop- 
ment of  cheap  ways  of  obtaining  electrical  energy  the 
problem  has  been  solved,  and  the  metal  is  now  produced 
by  the  electrolysis  of  aluminium  oxide  (A12O3)  dissolved  in 
melted  cryolite  — a  method  devised  by  the  American  chemist 


*FiG.  158.    Diagram  illustrating  the  manufacture  of  aluminium 

Hall  (Fig.  157)  in  1886.  The  annual  production  of  alumin- 
ium now  exceeds  50,000,000  Ib.  During  the  last  twenty-five 
years  its  cost  has  decreased  from  $5  to  20  cents  per  pound. 

Metallurgy.  An  iron  box  A  (Fig.  158)  about  8ft.  long  and 
6  ft.  wide  is  connected  with  a  powerful  electrical  generator  in 
such  a  way  as  to  serve  as  the  cathode  upon  which  the  aluminium 
is  deposited.  Three  or  four  rows  of  carbon  rods  7?,  B  dip  into  the 
box  and  serve  as  the  anodes.  The  box  is  partly  filled  with 
cryolite,  and  the  current  is  turned  on,  generating  enough  heat 
to  melt  the  cryolite.  Aluminium  oxide  is  then  added,  and  acts' 
as  an  electrolyte,  being  decomposed  into  aluminium  and  oxygen. 


324  FIRST  COURSE  IN  CHEMISTRY 

The  temperature  is  maintained  above  the  melting  point  of 
aluminium,  and  the  liquid  metal,  being  heavier  than  cryolite, 
sinks  to  the  bottom  of  the  vessel,  from  which  it  is  tapped  off 
from  time  to  time  through  the  tap  hole  C. 

Properties.  Aluminium  is  a  tin-white  metal  which  melts 
at  658.7°  and  is  very  light,  its  density  being  about  one 
third  that  of  iron.  It  is  stiff  and  strong,  and  with  frequent 
heating  can  be  rolled  into  thin  foil.  It  is  a  good  conductor 
of  heat  and  electricity,  though  not  so  good  as  copper  for  a 
given  cross  section  of  wire. 

Aluminium  is  not  perceptibly  acted  on  by  boiling  water, 
and  moist  air  merely  dims  its  luster.  Further  action  is  pre- 
vented in  each  case  by  the  formation  of  an  extremely  thin 
film  of  oxide  upon  the  surface  of  the  metal.  When  heated 
in  oxygen  it  burns  with  great  energy  and  with  the  liberation 
of  much  heat.  It  is  therefore  a  good  reducing  agent.  Hydro- 
chloric acid  acts  upon  it,  forming  aluminium  chloride  ;  nitric 
acid  and  dilute  sulfuric  acid  have  almost  no  action  on  it ; 
but  hot  concentrated  sulfuric  acid  acts  upon  it  in  the  same 
way  as  upon  copper,  forming  aluminium  sulfate.  Alkalies 
readily  attack  it,  liberating  hydrogen : 

2  Al  +  6  KOH *•  2  A1(OK)3  +  3  Ha 

Salt  solutions,  such  as  sea  water,  corrode  the  metal  rapidly. 
Uses  of  aluminium.  These  properties  suggest  many  uses 
for  the  metal.  Its  lightness,  strength,  and  inactivity  toward 
air  and  water  make  it  well  adapted  for  many  construction 
and  manufacturing  purposes.  These  same  properties  have 
led  to  its  extensive  use  in  the  manufacture  of  cooking  uten- 
sils. Owing  to  its  small  resistance  to  electrical  currents,  it 
is  replacing  copper  to  some  extent  in  electrical  construction, 
especially  for  trolley  and  power  wires.  In  the  form  of  a 
powder  suspended  in  a  suitable  liquid  it  makes  a  silvery 


ALUMINIUM 


325 


paint  used  to  cover  iron  pipes  and  lantern  curtains.  The 
greatest  use  of  aluminium  is  in  the  steel  industry  (p.  348). 
Aluminium  bronze,  consisting  of  about  90  per  cent  cop- 
per and  10  per  cent  aluminium,  has  a  pure-golden  color, 
is  strong  and  malleable,  is  easily  cast,  and  is  permanent  in 
the  air.  Magnalium  is  an  alloy  of  aluminium  and  magne- 
sium. It  is  light  and  rigid  and  is  used  for  balance  beams. 

Goldschmidt  reduction  process.  Al- 
uminium is  frequently  employed  as  a 
powerful  reducing  agent,  many  me- 
tallic oxides  which  resist  reduction  by 
carbon  being  readily  reduced  by  it. 
The  aluminium,  in  the  form  of  a  fine 
powder,  is  mixed  with  the  metallic 
oxide,  together  with  some  substance 
such  as  fluorite  to  act  as  a  flux.  The 
mixture  is  ignited,  and  the  aluminium 
unites  with  the  oxygen  of  the  me- 
tallic oxide,  liberating  the  metal. 
This  collects  in  a  fused  condition 
under  the  melted  fluorite. 


FIG.  159.  Welding  a  rail 
with  thermite 


Thermite  welding  process.  The  property  possessed  by  alu- 
minium of  reducing  oxides  with  the  liberation  of  a  large 
amount  of  heat  is  turned  into  practical  account  in  the  welding 
of  metals.  The  German  chemist  Goldschmidt  was  the  first  to 
use  aluminium  for  this  purpose.  The  welding  of  metals  by 
this  method  may  be  illustrated  by  a  single  example,  namely,  the 
welding  of  car  rails — a  process  often  carried  out  in  connection 
with  electric  railways  to  secure  good  electrical  connection.  The 
ends  of  the  rails  are  accurately  aligned  and  thoroughly  cleaned. 
A  sand  mold  A  (Fig.  159)  is  then  clamped  about  the  ends  of 
the  rail,  leaving  sufficient  space  so  that  the  metal  can  flow  in. 
The  ends  of  the  rails  are  heated  to  redness  by  the  flame  from 
a  gasoline  compressed-air  torch  directed  into  the  opening  in 


326  FIRST  COURSE  IN  CHEMISTRY 

the  mold.  Just  over  the  opening  is  placed  the  conical-shaped 
crucible  B,  which  contains  a  mixture  of  iron,  metallic  oxides, 
and  aluminium.  When  the  ends  of  the  rails  have  -been  heated 
to  redness  by  the  torch,  the  mixture  in  the  crucible  is  ignited, 
and  after  a  few  seconds  the  crucible  is  opened  at  the  bottom, 
and  the  molten  metal  resulting  from  the  reaction  in  the  cruci- 
ble is  allowed  to  flow  into  the  mold.  In  this  way  the  molten 
metal  surrounds  the  ends  of  the  rails  and,  as  it  cools,  welds 
them  firmly  together.  A  mixture  of  the  metallic  oxides  and 
aluminium  ready  for  use  in  welding  is  sold  under  the  name 
of  therm ite. 

Aluminium  oxide  (A1203).  This  substance  occurs  in  sev- 
eral forms  in  nature.  The  relatively  pure  crystals  are  called 
corundum ;  emery  is  a  variety  colored  dark  gray  or  black, 
usually  by  iron  compounds.  In  transparent  crystals,  tinted 
different  colors  by  traces  of  impurities,  it  forms  such  pre- 
cious stones  as  the  sapphire,  ruby,  topaz,  and  oriental 
amethyst.  All  these  varieties  are  very  hard,  falling  little 
short  of  the  diamond  in  this  respect.  The  cheaper  forms, 
corundum  and  emery,  are  used  for  cutting  and  grinding 
purposes.  Chemically  pure  aluminium  oxide  can  be  made 
by  igniting  the  hydroxide,  when  it  forms  a  white  powder : 

2  A1(OH)8 >•  Al/)3  +  3  H20 

The  artificially  prepared  oxide  is  largely  used  in  the  prep- 
aration of  aluminium.  Some  laboratory  utensils  such  as 
crucibles  and  tubes  are  made  of  aluminium  oxide,  which  is 
given  the  trade  name  Alundum.  The  same  material  is  used 
for  cutting  and  polishing  metals. 

Artificial  gems.  A  number  of  gems  are  now  prepared  in 
the  laboratory  from  molten  aluminium  oxide.  The  white  sap- 
phires so  extensively  advertised  are  simply  the  pure  oxide. 
By  incorporating  with  the  melted  oxide  small  percentages  of 


ALUMINIUM  327 

certain  metallic  oxides,  different  tints  or  colors  are  obtained, 
and  in  this  way  are  prepared  such  gems  as  the  ruby,  the 
oriental  amethyst,  and  the  yellow  and  blue  sapphires,  which 
are  practically  identical  in  composition  and  properties  with  the 
natural  stones. 

Aluminium  hydroxide  (A1(OH)3).  The  hydroxide  can  be 
prepared  by  adding  ammonium  hydroxide  to  any  soluble 
aluminium  salt,  forming  a  colloidal  precipitate  which  is 
insoluble  in  water  but  very  hard  to  filter.  When  heated,  it 
is  decomposed,  forming  the  oxide  and  water.  It  dissolves 
in  most  acids  to  form  soluble  salts,  and  in  the  strong  bases 
to  form  aluminates,  as  indicated  in  the  equations 

A1(QH)3  +  3  HC1  — >•  A1C18  +  3  H2O 
A1(OH)3  +  3  NaOH *-  Al(ONa)3  +  3  H2O 

It  may  act,  therefore,  either  as  a  weak  base  or  as  a 
weak  acid,  its  action  depending  upon  the  character  of  the 
substances  with  which  it  is  in  contact. 

Water  purification.  The  value  of  aluminium  hydroxide  in 
the  purification  of  water  (p.  44)  is  due  largely  to  its  colloidal 
or  gelatinous  character  when  freshly  formed  by  precipitation. 
After  being  stirred  through  the  water  it  is  allowed  to  slowly 
settle,  and  in  so  doing  it  carries  with  it  any  suspended  matter 
present,  including  microorganisms  and  coloring  materials.  In- 
stead of  adding  aluminium  hydroxide  itself  to  the  water,  it  is 
more  economical  and  effective  to  produce  it  by  precipitation. 
This  is  done  by  dissolving  in  the  water  some  cheap  salt  which 
readily  hydrolyzes,  such  as  aluminium  sulfate : 

A12(S04)3  +  6  H20 >•  2  A1(OH)8  +  3  H2S04 

There  is  always  sufficient  basic  material  present  in  the  water 
to  combine  with  the  sulfuric  acid  set  free,  so  that  no  acid  is 
left  in  the  water  as  a  result  of  this  treatment. 


328 


FIRST  COURSE  IN  CHEMISTRY 


Fig.  160  illustrates  the  use  of  aluminium  sulfate  in  purifying 
water.    The  cylinder  A  contains  impure  water.    B  is  a  similar 

cylinder   of   water   to  which 

'  •«•».. Mi:_jy»    e"trz^  <Z.I^p    I     some  aluminium  sulfate  has 

been  added.    The  aluminium 
hydroxide  formed  by  hydrol- 
ysis is  slowly  settling  in -the 
>MBMM«i^l      water,  carrying   with   it   the 

Mr 


A 


B 


impurities.  The  appearance 
of  the  water  after  settling  is 
shown  in  C. 


FIG.  100.    Purification  of  water  by 
aluminium  sulfate 


Alums.  Aluminium  sul- 
fate can  be  prepared  by  the 
action  of  sulfuric  acid  upon 
the  mineral  bauxite.  It  has 
the  property  of  combining 
with  the  sulfates  of  the 
alkali  metals  to  form  com- 
pounds called  alums.  Thus,  with  potassium  sulfate  the 
reaction  is  expressed  by  the  equation 

K2S04  +  A12(S04)3  +  24  H20  — >-  2  (KA1(SO4)2  -  12  H2O) 

The  sulfates  of  other  trivalent  metals  can  form  similar  com- 
pounds with  the  alkali  sulfates,  and  these  compounds  are  also 
called  alums,  though  they  contain  no  aluminium.  They  all  crys- 
tallize in  octahedra  and  contain  12  molecules  of  water  of  hydra- 
tion.  The  alums  most  frequently  prepared  are  the  following : 

Potassium  alum KA1(SO4)2  •  12  II2O 

Ammonium  alum NH4A1(SO4),  •  12  H2O 

Ammonium  iron  alum     ....  NH4Fe(SO4)2  .  12  II2O 

Potassium  chrome  alum  ....  KCr(SO4)2  •  12  II2O 

Very  large,  well-formed  crystals  of  an  alum  can  be  prepared 
by  suspending  a  small  crystal  by  a  thread  in  a  saturated  solu- 
tion of  the  alum,  as  shown  in  Fig.  161.  The  small  crystal 
slowly  grows  and  assumes  a  very  perfect  form. 


ALUMINIUM 


329 


Hydrolysis  of  salts  of  aluminium.  While  aluminium 
hydroxide  forms  fairly  stable  salts  with  strong  acids,  it  is 
such  a  weak  base  that  its  salts  with  weak  acids  are  readily 
hydrolyzed  (p.  138).  Thus,  when  an  aluminium  salt  and 
a  soluble  carbonate  are  brought  together  in  solution,  we 
should  expect  to  have  aluminium  carbonate  precipitated 
according  to  the  equation 

3  Na2C03  +  2  A1C13  -  >-  A12(CO3)3  +  6  NaCl 

But  if  it  is  formed  at  all,  it  instantly  begins  to  hydrolyze, 
the  products  of  the  hydrolysis  being  aluminium  hydroxide 
and  carbonic  acid  : 


A12(C03) 


6  H20 


A1(OH) 


3  H2CO3 


Aerating  agents  used  in  baking.  In  preparing  foods  made 
largely  from  dough,  such  as  bread,  biscuits,  and  cake,  it  is 
essential  that  some  aerating  agent  be  used 
to  render  the  food  light  and  wholesome. 
The  aerating  agent  used  in  all  cases  is 
carbon  dioxide.  This  is  generated  in  the 
dough  and,  pushing  its  way  through  the 
mass,  renders  it  porous  and  light.  The  fol- 
lowing methods  are  used  for  generating 
the  gas  in  baking  : 


FIG.  161.  The  for- 
mation of  a  crystal 
of  alum 


1.  By  alcoholic  fermentation.    As  we  have 
seen,  this  is  the  method  generally  used  in 
making  bread  (p.  231). 

2.  By  the  action  of  sour  milk  on  sodium  bi- 
carbonate.   The  lactic  acid  present  in  the  sour  milk  (p.  222) 
slowly  acts  upon  the  bicarbonate,  liberating  carbon  dioxide : 

H  -  C.H.O,  +  NaHC03  — »•  NaC^.O.  +  H20  +  C02 

This  method  has  largely  been  replaced  by  the  following  one : 


330  FIRST  COURSE  IN  CHEMISTRY 

3.  By  the  action  of  an  acid  salt  or  alum  upon  sodium  bicarbon- 
ate; baking  powders.  Mixtures  of  sodium  bicarbonate,  flour  (or 
starch),  and  some  substance  that  will  act  upon  the  bicarbon- 
ate to  liberate  carbon  dioxide  are  known  as  baking  powders. 
The  compounds  commonly  employed  for  liberating  the  carbon 
dioxide  are  either  alum,  cream  of  tartar  (potassium  bitartrate) 
(p.  237),  or  calcium  acid  phosphate,  and  baking  powders  are 
known  as  alum  baking-powders,  cream  of  tartar  baking  powders, 
OY  phosphate  baking  powders  ^  according  to  whether  they  contain 
the  one  or  other  of  these  constituents.  The  reactions  take  place 
only  in  the  presence  of  water  ;  hence  the  use  of  the  flour,  which, 
by  absorbing  any  moisture  that  may  be  present,  prevents  the 
powder  from  losing  its  strength  until  used.  In  place  of  alum 
a  mixture  of  sodium  sulfate  and  aluminium  sulfate  known  as 
cream  of  tartar  substitute,  or  simply  as  C.  T.  S.,  is  now  being 
largely  used.  Alum  baking  powders  are  much  cheaper  than 
those  made  from  cream  of  tartar. 

The  reactions  of  baking  powders.  The  reactions  that  take  place 
when  water  is  added  to  each  of  the  classes  of  baking  powders 
are  represented  in  the  following  equations  : 

Alum  (supposing  that  the  alum  present  is  potassium  alum)  : 

2  KA1(S04)2  +  6  NaHC03  -  >- 

2  A1(OH)8  +  3  NaaS04  +  K2S04  +  6  C0a 
Cream  of  tartar  : 


KHC4H406  +  NaHCOs  -  *•  KNaC4H406  -f-  H20  +  C02 
Phosphate  : 
CaH4(P04)2  +  2  NaHC03  -  >• 


CaHP04  +  Na2HP04  +  2  H20  +  2  C02 


Dyes  and  dyeing.  To  gain  an  understanding  of  the  art 
of  dyeing  it  is  necessary  to  keep  in  mind  (1)  the  charac- 
teristics that  a  good  dye  must  have,  and  (2)  the  process 
of  fastening  the  dye  upon  the  fabric. 


DYES  AND  DYEING  331 

1.  The  dyes.   The  requisites  of  a  good  dye  are  as  follows : 
(a)  it  must  have  an  acceptable  color ;   (/>)  it  must  not  in- 
jure the  fibers ;   (c)  it  must  dye  fast ;  in  other  words,  the 
cloth  after  having  been  dyed  must  retain  its  color  when 
washed  with  water;  (d)  it  must  not  fade  too  easily.    In 
olden  times  nearly  all  the  dyes  used  were  extracted  from 
plants  and  trees.    These  dyes  give  rather  dull  but  pleasing 
colors,  and  the  beautiful  tones  of  oriental  rugs  and  tapestries 
are  primarily  due  to  their  use.    In  1856  the  Englishman, 
W.  H.  Perkin,  then  but  eighteen  years  of  age,  prepared  the 
first  aniline  dye.    The  preparation  of  others  soon  followed, 
until  to-day  thousands  of  these  dyes  are  manufactured,  of 
every  variety  of  color,  while  new  ones  are  constantly  being 
added  to  the  list.    Moreover,  two  of  our  most  common  dyes, 
indigo  and  alizarin,   which  were  formerly  obtained  from 
vegetable  sources,  are  now  manufactured  (p.  232)  from 
coal-tar  hydrocarbons.  These  aniline  dyes  (or  coal-tar  dyes) 
have  a  much  higher  coloring  power  than  the  vegetable  dyes, 
and  have  almost  entirely  superseded  them.    They  are  all 
very  complex  compounds. 

2.  The  process  of  dyeing.    This  process  consists  in  fixing 
the   dye   uniformly   upon   the   fabric.    The   animal   fibers, 
namely,  wool  and  silk,  are  as  a  rule  more  readily  dyed  than 
cotton,  which  is  a  vegetable  fiber.    To  dye  wool  and  silk 
with  most  dyes  it  is  only  necessary  to  steep  the  fabric  in 
a  solution  of  the  dye.    Cotton  fabrics,  when  treated  in  this 
way,  will  become  colored,  but  as  a  rule  the  color  is  not  fast. 
Cotton  fabrics  may  be  dyed  fast  in  the  following  way : 

The  cloth  is  first  soaked  in  a  solution  of  an  aluminium  salt 
(or  a  similar  substance),  which  readily  undergoes  hydrolysis. 
The  cloth  is  then  exposed  to  the  action  of  steam,  which 
decomposes  the  salt,  leaving  the  hydroxide  thoroughly  in- 
corporated in  the  fiber.  The  cloth  is  then  steeped  in  the 


332  FIRST  COURSE  IN  CHEMISTRY 

dye,  which  is  absorbed  by  the  aluminium  hydroxide,  and 
is  in  consequence  fastened,  or  "  fixed,"  upon  the  fiber. 
Aluminium  hydroxide  and  other  substances  which  act  in 
the  same  way  are  called  mordants.  The  same  dye  will 
often  give  different  colors  with  different  mordants. 

Lakes.  The  compounds  which  serve  well  as  mordants  may 
be  precipitated  in  solutions  containing  various  dyes,  and 
the  precipitate  will  be  highly  colored,  though  not  always 
of  the  same  color  as  the  dye.  Colored  precipitates  of  this 
kind  are  called  lakes  and  are  used  as  pigments  in  paints. 

EXERCISES 

1.  Why   should  not   aluminium   be   made  by  electrolyzing  the 
melted  hydroxide  as  with  so  many  other  metals  ? 

2.  What  familiar  articles  can  you  mention  which  are  now  made 
of  aluminium? 

3.  What  is  a  colloidal  substance? 

4.  Why  do  the  directions  for  using  aluminium  cooking  utensils 
state  that   such  utensils  must  not  be  washed  in  strongly  alkaline 
solutions  ? 

5.  What  is  the  meaning  of  the  word  mordant  (see  dictionary)? 

6.  Calculate  the  weights  of  substances  necessary  for  the  prepara- 
tion of  10  Ib.  of  cream  of  tartar  baking  powder,  supposing  that  such 
powders  contain  25  per  cent  of  starch.    A  ns.  Starch,  40  oz. ;  bicar- 
bonate of  soda,  37  oz. ;  cream  of  tartar,  83  oz. 

7.  What  volume  of  carbon  dioxide  measured  at  200°  would  be 
evolved  by  6  g.   (approximately  2  level  teaspoonfuls)  of  a  baking 
powder  of  the  composition  found  in  problem  9?    Ans.  849  cc. 

8.  It  will  be  of  interest  to  calculate  the  data  in  problems  6  and  7  for 
each  of  the  other  classes  of  baking  powders  and  to  compare  the  results. 

TOPICS  FOR  THEMES 

The  dyeing  of  cloth  (Lassar-Cohn,  Chemistry  in  Daily  Life). 

The  advantages  of  aluminium  for  cooking  utensils.  (Consult 
dealers.) 

Artificial  gems  (Duncan,  The  Chemistry  of  Commerce). 


CHAPTER  XXXV 

ALUMINIUM  SILICATES  AND  THEIR  COMMERCIAL 
APPLICATIONS 

Aluminium  silicates.  One  of  the  most  common  constit- 
uents of  rocks  is  feldspar  (KAlSi3O8),  a  mixed  salt  of 
potassium  and  aluminium  with  polysilicic  acid  (H4Si3O8). 
Under  the  influence  of  moisture,  carbon  dioxide,  and 
changes-  of  temperature  this  substance  is  constantly  being 
broken  down  into  soluble  potassium  compounds  and  alu- 
minium silicate  (Al2Si2O?  -  2  H2O).  In  relatively  pure  con- 
dition it  is  called  kaolin  and  is  a  soft,  plastic  mineral ;  in 
the  impure  state,  mixed  with  sand  and  other  substances, 
it  forms  common  clay.  Mica  is  another  very  abundant  min- 
eral, having  a  varying  composition,  but  being  essentially  of 
the  formula  KAlSiO4.  Serpentine,  talc,  asbestos,  and  meer- 
schaum are  important  complex  silicates  of  aluminium  and 
magnesium;  granite  is  a  mechanical  mixture  of  quartz, 
feldspar,  and  mica  and  is  therefore  rich  in  aluminium. 
Fuller  s  earth  is  a  peculiar  form  of  aluminium  silicate, 
which  is  used  as  a  filtering  material  for  decolorizing  oils, 
especially  cottonseed  oil. 

Clay  products.  The  crudest  forms  of  clay  products,  such 
as  porous  brick  and  draintile,  have  little  chemistry  involved 
in  their  manufacture.  Natural  clay  is  molded  into  the  re- 
quired form,  dried,  and  then  burned  in  a  kiln,  but  not  to  a 
temperature  at  which  the  materials  soften.  In  this  process 
the  nearly  colorless  iron  compounds  in  the  clay  are  con- 
verted into  colored  compounds  which  give  the  usual  red 

333 


334  FIRST  COURSE  IN  CHEMISTRY 

color  to  these  articles.  In  making  vitrified  brick  the  tem- 
perature is  raised  to  the  point  at  which  fusion  begins,  so 
that  the  brick  is  partially  changed  to  a  kind  of  glass. 

White  pottery.  This  term  is  applied  to  a  variety  of 
articles  varying  from  the  crudest  porcelain  to  the  finest 
chinaware.  While  the  processes  used  in  the  manufacture 
of  the  articles  differ  in  details,  fundamentally  they  are  the 
same  and  may  be  described  under  three  heads :  namely, 

(1)  the  preparation 
of  the  body  of  the 
ware,  (2)  the  proc- 
ess of  glazing,  and 
(3)  the  decoration. 

1.  The  body  of  the 
ware.  The  materials 
used  consist  of  an  ar- 
tificially compounded 
clay  made  from  kaolin, 
plastic  clay,  and  pul- 
verized feldspar.  This 

FIG.  162.  The  manufacture  of  pottery  ;  mold-     mixture  is  plastic  and 
ing  the  plastic  material  into  form  is  worked  into  the  de- 

sired shape  by  molds 

or  on  a  potter's  wheel  (Fig.  162).  The  ware  is  then  dried  and 
burned  in  a  kiln  (Fig.  163)  until  vitrified,  and  in  this  form  is 
known  as  bisque.  This  is  usually  porous  and  must  therefore 
be  glazed  to  render  it  nonabsorbent. 

2.  The  glaze.  The  glaze  is  a  fusible  glass  which  is  melted 
over  the  surface  of  the  body.  The  constituents  of  the  glaze 
are  quartz,  feldspar,  and  various  metallic  oxides,  often  mixed 
with  a  little  boric  oxide.  These  materials  are  finely  ground  and 
mixed  with  water  to  a  paste.  Sometimes  they  are  first  fused 
into  a  glass,  which  is  then  powdered  and  made  into  the  paste. 
The  bisque  is  dipped  into  the  glaze  paste,  dried,  and  fired 
until  the  glaze  materials  melt  and  flow  evenly  over  the  surface. 


ALUMINIUM  SILICATES 


335 


3.  The  decoration.  If  the  article  is  to  be  decorated,  the 
design  may  be  painted  upon  the  body  before  glazing,  or  it 
may  be  painted  upon  the  glaze  and  the  article  fired  again, 
the  pigments  melting  into  the  glaze.  In  the  former  case 
the  pigments  used  are  as  a  rule  metallic  oxides  of  various 
colors,  while  in  the  latter  case  they  are  often  colored  glasses. 

Cement.  The  term 
cement  as  ordina- 
rily used  at  present 
is  applied  to  those 
mortars  which  pos- 
sess the  property  of 
hardening  in  water 
as  well  as  in  air. 
These  cements  are 
silicate  bodies,  usually 
very  highly  basic  in 
character,  and  when 
ground  fine  and 
mixed  with  water, 
they  undergo  com- 
plex reactionsresult-  FlG' 163'  The  manufacture  °f  P°ttery  J  burn' 

mg  the  ware  in  a  kiln 
ing  in  the  formation 

of  a  hard,  rocklike  mass.  A  number  of  different  classes 
of  cements  are  known,  the  most  important  of  which  is 
called  Portland  cement. 

Composition  of  Portland  cement.  The  essential  ingredients 
of  Portland  cement,  together  with  the  general  limits  of  each 
ingredient,  are  as  follows  : 


SiO, 19  to  26% 


ALO, 


4  to  11% 


MgO Oto5% 

SO 0  to  2.5% 


Fe2O3 2  to  5% 

CaO 58  to  67% 


336  FIKST  COUKSE  IX  CHEMISTRY 

Manufacture  of  Portland  cement.  The  materials  most  com- 
monly employed  are  limestone  or  marl  and  clay  or  shale.  In 
general,  however,  any  substance  may  be  used  which  furnishes 
the  ingredients  listed  in  the  above  table.  Among  the  sub- 
stances so  used  is  blast-furnace  slag,  which  is  an  impure 
calcium-aluminium  silicate. 

The  materials  to  be  used  are  coarsely  ground  and  then 
mixed  together  in  the  proper  proportions  and  finely  pulver- 
ized. The  resulting  mixture  is  run  into  a  furnace  and  burned 
to  a  temperature  just  short  of  fusion,  at  which  temperature  it 


FIG.  164.    A  bridge  built  of  reinforced  concrete 

vitrifies,  forming  a  grayish  mass  known  as  clinker.  Finally, 
the  clinker  is  ground  to  a  fine  powder.  Gypsum  is  often  added 
in  the  process ;  this  acts  as  a  negative  catalyzer,  retarding  the 
hardening,  or  setting,  of  the  cement. 

The  setting  of  cement.  The  reactions  which  take  place 
upon  the  addition  of  water  to  cement,  and  which  result  in 
the  formation  of  a  hard,  rocklike  mass,  are  not  thoroughly 
understood.  The  constituents  of  the  cement  apparently 
undergo  hydrolysis  when  they  come  in  contact  with  water. 
The  resulting  compounds  unite  with  water  producing  hy- 
drates. These  hydrates  are  crystalline  in  character  and 
form  a  hard,  compact  mass. 


ALUMINIUM  SILICATES  337 

Growing  importance  of  cement.  Cement  is  rapidly  coming  into 
use  for  a  great  variety  of  purposes.  It  is  often  used  in  place 
of  mortar  in  the  construction  of  brick  buildings.  Mixed  with 
crushed  stone  and  sand  it  forms  concrete,  which  is  used  in 
foundation  work  for  buildings  and  street  paving.  It  is  also 
used  in  making  artificial  stone,  terra  cotta  trimmings  for  build- 
ings, artificial-stone  walks  and  floors,  fence  posts,  and  the  like. 
It  is  being  used  more  and  more  for  making  articles  which  were 
formerly  made  of  wood  or  stone,  and  the  entire  walls  of  build- 
ings are  sometimes  made  of  cement  blocks  or  of  concrete.  Iron 
rods  or  wire  are  often  embedded  in  the  concrete  before  it  sets, 
to  give  it  greater  strength,  and  this  is  called  reinforced  concrete. 

EXERCISES 

1.  In  the  manufacture  of  pottery  why  is  the  glaze  made  more 
fusible  than  the  body  of  the  ware? 

2.  Suppose  that  the  glaze  and  the  body  expand  and  contract  at 
different  rates  with  changes  in  temperatures,  what  will  be  the  result  ? 

3.  What  is  the  meaning  of  the  word  vitrify! 

4.  What  is  a  catalyzer?    What  is  a  negative  catalyzer? 

5.  Why  can  cement  be  used  as  mortar  in  colder  weather  than 
ordinary  mortar  ? 

6.  What  is  a  polysilicic  acid? 

7.  What  weight  of  kaolin  will  result  from  the  weathering  of 
1  ton  of  feldspar?    Ans.  927  Ib. 

TOPICS  FOR  THEMES 

The  making  of  porcelain  (Lassar-Cohn,  Chemistry  in  Daily  Life). 

How  cement  is  made  (Rogers  and  Aubert,  Industrial  Chemistry). 

The  history  of  a  china  dish  (Rogers  and  Aubert,  Industrial 
Chemistry). 


CHAPTER  XXXVI 
THE   IRON  FAMILY 


NAME 

SYMBOL 

ATOMIC 
WEIGHT 

DEN.SITY 

MELT  INC; 
POINT 

OXIDES 

Iron      .     .     . 

Fe 

55.84 

7.86 

1530° 

FeO,  Fe2O3 

Cobalt  .     .     . 

Co 

58.97 

8.6 

1478° 

CoO,  Co2O3 

Nickel  .     .     . 

Ni 

58.68 

8.9 

1452° 

NiO,  Ni2O3 

The  family.  The  elements  iron,  cobalt,  and  nickel  form 
a  group  in  the  eighth  column  of  the  periodic  table.  The 
atomic  weights  of  the  three  are  very  close  together,  and 
their  properties  are  very  much  alike. 

IRON 

Occurrence.  The  element  iron  has  long  been  known,  since 
its  ores  are  very  abundant  and  it  is  not  difficult  to  prepare 
the  metal  from  them  in  fairly  pure  condition.  It  occurs  in 
large  deposits  as  oxides,  sulfides,  and  carbonates,  and  in 
smaller  quantities  in  a  great  variety  of  minerals.  Indeed, 
very  few  rocks  or  soils  are  free  from  small  percentages  of 
iron.  It  is  a  constituent  of  the  chlorophyll  of  plants  and 
the  haemoglobin  of  the  blood  of  animals,  and  therefore  plays 
an  important  part  in  life  processes.  Many  meteorites  are 
largely  iron,  usually  alloyed  with  a  little  nickel. 

Pure  iron.  Pure  iron  may  be  prepared  by  the  electrolysis 
of  a  solution  of  iron  sulfate  between  iron  electrodes,  though 
it  is  difficult  to  free  it  entirely  from  hydrogen  in  this  way. 

338 


THE  IRON  FAMILY  339 

It  is  prepared  in  practically  pure  condition  by  the  open- 
hearth  method  (p.  346).  It  is  a  silvery  metal  which  melts 
at  1530°.  It  is  ductile  and  malleable  and  almost  as  soft 
as  aluminium.  It  is  especially  well  adapted  to  the  manufac- 
ture of  electromagnets,  since  it  acquires  and  loses  magnetic 
properties  more  readily  than  do  the  ordinary  varieties  of 
iron.  It  is  also  used  for  purposes  where  resistance  to  corro- 
sion is  desired,  for  it  does  not  rust  rapidly. 

The  iron  of  commerce.  Iron  differs  from  most  of  the 
other  metals  used  in  the  industries  in  that  the  pure  metal 
is  seldom  obtained  and  is  of  limited  application,  while  that 
containing  small  percentages  of  other  elements  exhibits  a 
wide  variety  of  properties  which  make  it  of  the  greatest 
value  for  many  different  purposes. 

Carbon  is  always  present  in  amounts  which  vary  from  a 
mere  trace  to  about  7  per  cent.  According  to  the  condition 
of  treatment,  the  carbon  may  be  in  the  form  of  graphite 
scattered  through  the  iron,  or  it  may  occur  as  a  solid  solution 
of  carbon  in  iron,  or  as  carbides  of  iron,  one  of  the  most 
important  of  which  has  the  formula  FegC,  and  is  called 
cementite.  Manganese,  silicon,  and  traces  of  phosphorus  and 
sulfur,  together  with  a  little  oxygen,  are  also  present. 

The  properties  of  the  iron  are  so  much  modified  by  the 
percentages  of  these  elements,  by  their  form  of  combination, 
and  by  the  treatment  of  the  metal  during  its  production, 
that  many  varieties  of  iron  are  recognized  in  commerce, 
the  chief  of  which  are  cast  iron,  wrought  iron,  and  steel. 

Materials  used  in  metallurgy  of  iron.  Four  different 
classes  of  materials  are  used  in  the  metallurgy  of  iron : 

1.  Iron  ore.  The  ores  most  frequently  employed  are  the 
following : 

Hematite Fe2O3      Siderite  .     .     .     FeCO3 

Magnetite Fe3O4      Limonite     .     .     2  Fe2O3  •  3  H2O 


340 


FIKST  COURSE  IN  CHEMLSTBY 


While  iron  ore  is  mined  in  a  number  of  different  localities 
in  the  United  States,  the  great  center  of  production  is  in 
the  neighborhood  of  Lake  Superior,  the  ore  being  chiefly 
hematite.  Large  amounts  are  also  mined  near  Birmingham, 
Alabama.  Fig.  165  represents  one  of  the  large  mines  in 
Minnesota. 

2.  Carbon.  Carbon  in  some  form  is  necessary  both  as  a 
fuel  and  as  a  reducing  agent.  In  former  times  wood  charcoal 

• =====aaa=ssamsssa!S!S=a^     was  use(l  to  supply 

the  carbon,  but  now 
coke  is  almost  uni- 
versally used. 

3.  Hot  air.  To 
maintain  the  high 
temperature  required 
for  the  reduction  of 
iron,  a  very  active 
combustion  of  fuel 
is  necessary.  This  is 
secured  by  forcing  a 

strong  blast  of  hot  air  into  the  lower  part  of  the  furnace 
during  the  reduction  process. 

4.  Flux.  All  the  materials  which  enter  the  furnace  must 
leave  it  again,  either  in  the  form  of  gases  or  as  liquids. 
The  iron  is  drawn  off  as  the  liquid  metal  after  its  reduction, 
the  oxygen  with  which  it  was  combined  escaping  as  oxides 
of  carbon.  To  secure  the  removal  of  the  earthy  matter 
charged  into  the  furnace  along  with  the  ore,  materials  are 
added  to  the  charge  which  will  combine  with  the  impurities 
in  the  ore,  forming  a  liquid.  The  material  added  for  this 
purpose  is  called  the  flux  and  usually  consists  of  limestone. 
The  liquid  produced  from  the  flux  and  the  ore  is  called 
dag.  It  is  a  variety  of  readily  fusible  glass. 


FIG.  1G5.    Mining  iron  ore  in  Minnesota 


THE  IRON  FAMILY 


341 


Cast  iron.  Ordinarily  the  first  step  in  the  manufacture 
of  any  variety  of  commercial  iron  is  the  production  of  cast 
iron.  The  ores  are  mixed  with  a  suitable  flux,  and  are 
reduced  by  heating  with  coke. 

Blast-furnace  process.  The  reduction  is  carried  out  in  a  large 
tower,  called  a  blast  furnace  (Fig.  166).  This  is  usually  80  ft. 
high  and  20  ft.  in  internal  diameter  at 
its  widest  part,  narrowing  somewhat 
toward  both  the  top  and  the  bottom. 
The  walls  are  built  of  steel  and  are 
lined  with  fire  brick.  The  base  is  pro- 
vided with  a  number  of  pipes  ,4,  called 
tuyeres,  through  which  hot  air  is  forced 
into  the  furnace.  The  tuyeres  are  sup- 
plied from  a  large  pipe  B,  which  girdles 
the  furnace.  At  the  base  of  the  fur- 
nace is  an  opening,  through  which  the 
liquid  metal  can  be  drawn  off  from 
time  to  time.  There  is  also  a  second 
opening  C,  somewhat  above  the  first, 
through  which  the  excess  of  slag  over- 
flows. The  top  is  closed  by  a  movable 
trap  Z>,  called  the  cone,  and  through 
this  the  materials  to  be  used  are  intro- 
duced.. The  gases  resulting  from  the 
combustion  of  the  fuel  and  the  reduc- 
tion of  the  ore,  together  with  the  ni- 
trogen of  the  air  admitted  through 
the  tuyeres,  escape  through  pipes  E. 
These  gases  are  very  hot  and  contain 

a  sufficient  percentage  of  carbon  monoxide  to  render  them 
combustible ;  they  are  accordingly  utilized  for  heating  the 
blast  of  air  admitted. through  the  tuyeres,  and  as  fuel  for  the 
engines. 

Charges  consisting  of  coke,  ore,  and  flux  in  proper  propor- 
tion are  at  intervals  introduced  into  the  furnace  through  the 


FIG.  166.    Diagram  of  a 
blast  furnace 


342 


FIRST  COUESE  IN  CHEMISTRY 


cone.  The  coke  burns  fiercely  in  the  hot-air  blast,  forming  car- 
bon dioxide,  Avhich  is  at  once  reduced  to  carbon  monoxide  as  it 
passes  over  the  highly  heated  carbon. 

Reduction  of  the  ore  begins  at  the  top  of  the  furnace  through 
the  action  of  the  carbon  monoxide.  As  the  ore  slowly  descends, 
the  reduction  is  completed,  and  the  resulting  iron  melts  and 
collects  as  a  liquid  in  the  bottom  of  the  furnace,  the  lighter 
slag  floating  above  it.  After  a  considerable  quantity  of  iron  has 


FIG.  167.   A  typical  plant  for  the  manufacture  of  cast  iron 

collected,  the  slag  is  drawn  off  through  C,  and  the  iron  is  run 
into  ladles  and  taken  to  the  converters  for  the  manufacture  of 
steel ;  or  it  is  run  into  sand  molds  and  cast  into  ingots  called 
pigs.  Fig.  168  shows  the  method  of  drawing  off  the  iron.  A 
small  hole  is  made  near  the  bottom  of  the  furnace,  and  the 
molten  iron  flows  out  and  down  the  central  channel  A  and 
into  the  sand  molds  along  the  sides  B,  where  it  solidifies. 

In  practice,  a  number  of  furnaces  are  usually  operated  to- 
gether, as  illustrated  in  Fig.  167,  which  shows  an  exterior 
view  of  a  modern  plant  for  making  cast  iron. 


THE  IRON  FAMILY  343 

Properties  of  cast  iron.  The  iron  produced  in  the  blast 
furnace  is  called  cast  iron.  It  varies  considerably  in  com- 
position, but  always  contains  over  2  per  cent  of  carbon, 
variable  amounts  of  silicon,  and,  at  least,  traces  of  phospho- 
rus and  sulfur.  The  form  in  which  the  carbon  is  present, 
whether  free  or  combined,  also  greatly  modifies  the  proper- 
ties of  the  iron.  In  general,  cast  iron  is  hard  and  brittle, 
and  melts  at  about  1100°.  It  cannot  be  welded  or  forged, 


FIG.  168.    Casting  pig  iron  from  a  blast  furnace 

but  is  easily  cast  in  sand  molds.  It  is  rigid,  but  not  elastic, 
and  its  tensile  strength  is  small.  It  is  used  for  making 
castings,  but  chiefly  as  a  starting  point  in  the  manufacture 
of  other  varieties  of  iron. 

Wrought  iron.  Wrought  iron  is  made  from  cast  iron  by 
burning  out  most  of  the  carbon,  silicon,  phosphorus,  and 
sulfur,  the  operation  being  conducted  in  what  is  called  a 
puddling  furnace. 

Wrought  iron  is  soft,  malleable,  and  ductile.  Its  tensile 
strength  is  greater  than  that  of  cast  iron,  but  less  than  that 
of  most  steel.  Its  melting  point  is  much  higher  than  that 


344 


FIRST  COURSE 


CHEMISTEY 


of  cast  iron,  and  if  melted,  it  is  changed  into  steel.  It  is  no 
longer  produced  to  the  same  relative  extent  as  in  former 
years,  since  soft  steel  can  be  made  at  a  less  cost  and  has 
almost  the  same  properties. 

Steel.  Steel,  like  wrought  iron,  is  made  from  cast  iron 
by  burning  out  a  part  of  the  carbon,  silicon,  phosphorus, 
and  sulfur,  but  the  processes  used  are  quite  different  from 
that  employed  in  the  manufacture  of  wrought  iron.  Nearly 
all  the  steel  of  commerce  produced  in  the  United  States  is 
made  by  one  of  two  general  methods 
known  as  the  acid  Bessemer  process  and 
the  basic  open-hearth  process. 

Acid  Bessemer  process.  In  the  acid 
Bessemer  process  the  furnaces  used 
are  lined  with  silica,  which,  it  will  be 
recalled,  is  an  acid  anhydride.  These 
furnaces  remove  from  the  cast  iron  the 
carbon  and  silicon,  but  not  the  phos- 
phorus and  sulfur.  The  process  is  there- 
fore employed  when  the  cast  iron  to  be 
used  is  low  in  phosphorus  and  sulfur. 


FIG.  169.    Diagram  of 
a  Bessemer  converter 


Details  of  operation  of  Bessemer  process.  This  process,  in- 
vented about  1880,  is  carried  out  in  great  egg-shaped  crucibles 
called  converters  (Fig.  169),  each  one  of  which  will  hold  as 
much  as  15  tons  of  steel.  The  converter  is  built  of  steel  and 
lined  with  silica.  It  is  mounted  on  trunnions,  so  that  it  can 
be  tipped  over  on  its  side  for  filling  and  emptying.  One  of  the 
trunnions  is  hollow,  and  a  pipe  connects  it  with  an  air  cham- 
ber J,  which  forms  a  false  bottom  to  the  converter.  The  true 
bottom  is  perforated,  so  that  air  can  be  forced  in  by  an  air 
blast  admitted  through  the  trunnion  and  the  air  chamber. 

White-hot  liquid  cast  iron  from  a  blast  furnace  is  run  into 
the  converter  through  its  open,  necklike  top  B,  the  converter 
being  tipped  over  to  receive  it ;  the  air  blast  is  then  turned  on, 


THE  IRON  FAMILY 


345 


and  the  converter  turned  to  a  nearly  vertical  position.  The 
carbon  and  silicon  in  the  iron  are  rapidly  oxidized  (first  the 
silicon  and  then  the  carbon),  the  oxidation  being  attended  by 
a  brilliant  flame  (Fig.  170).  The  heat  of  the  reaction,  largely 
due  to  the  combustion  of  silicon,  keeps  the  iron  in  a  molten 
condition.  The  air  blast  is  continued  until  the  character  of  the 
flame  shows  that  all  the  carbon  has  been  burned  away.  The 
process  requires  from  fifteen  to 
twenty  minutes,  and  when  it 
is  complete,  the  desired  quan- 
tity of  carbon  (generally  in  the 
form  of  high  carbon  iron  alloy) 
is  added  and  allowed  to  mix 
thoroughly  with  the  fluid.  The 
converter  is  then  tilted,  and 
the  steel  run  into  molds,  and 
the  ingots  so  formed  are  ham- 
mered or  rolled  into  rails  or 
other  objects. 

Basic  open-hearth  process. 
In  the  basic  open-hearth  proc- 
ess the  lining  of  the  furnace 
is  made  of  limestone  or  dolo- 
mite, both  of  which  act  as 
bases.  In  such  furnaces  the 
phosphorus  and  sulfur  are 
both  removed,  as  well  as  the  silicon  and  carbon.  The  pres- 
ence of  more  than  traces  of  phosphorus  and  sulfur  in  the 
finished  steel  renders  the  metal  so  brittle  that  it  is  worth- 
less. The  open-hearth  process,  therefore,  possesses  a  great 
advantage  over  the  acid  Bessemer  process  in  that  it  makes 
it  possible  to  utilize  iron  ores  (or  cast  iron  obtained  from 
them)  that  contain  appreciable  quantities  of  phosphorus  and 
sulfur.  The  operation  does  not  need  to  be  hastened,  and 
steel  of  any  desired  composition  can  be  produced. 


FIG.  170.    A  Bessemer  converter 
in  operation 


346 


FIKST   COURSE  IX  CHEMISTRY 


Details  of  the  open-hearth  process.  Fig.  171  shows  the  simpler 
parts  of  the  type  of  furnace  used  in  this  process.  The  hearth 
of  the  furnace  is  about  40  ft.  in  length,  12  ft.  in  width,  and  2  ft. 
in  depth,  and  is  lined  with  limestone  or  dolomite  (.1,  A).  Either 
gas  or  sprayed  oil  is  used  as  fuel.  Below  the  furnace  is  placed 
a  checker  work  of  brick  so  arranged  that  the  hot  products  of  com- 
bustion escaping  from  the  furnace  may  be  conducted  through 
it,  thus  heating  the  bricks  to  a  high  temperature.  Both  the  air 
necessary  for  combustion  and  the  gaseous  fuel  (unless  decom- 
posed by  heating,  as  in  the  case  of  natural  gas  and  sprayed  oil) 
are  preheated  by  passing  them  over  the  hot  bricks,  so  that  the 
temperature  reached  during  combustion  is  greatly  increased. 


Hot  Aii 

FIG.  171.    Diagram  of  an  open-hearth  furnace 


The  gas  entering  through  C  comes  in  contact  at  D  with  the 
hot  air  entering  through  /;,  and  a  vigorous  combustion  results, 
the  flame  passing  above  and  over  the  cast  iron  and  lime  with 
which  the  furnace  is  charged.  The  products  of  combustion  es- 
cape through  E  and  F.  At  the  temperature  reached,  the  carbon 
in  the  cast  iron  is  removed  in  the  form  of  the  oxide,  the  escap- 
ing gas  giving  the  melted  metal  the  appearance  of  boiling.  The 
silicon,  phosphorus,  and  sulfur  unite  with  oxygen  to  form  acid 
anhydrides ;  these  combine  with  the  lime  to  form  a  slag,  and  this 
rises  to  the  surface  of  the  melted  charge  and  is  easily  removed. 

When  a  test  shows  that  the  desired  percentage  of  carbon  is 
present,  the  melted  steel  is  run  into  large  ladles  and  then  into 
molds.  An  average  furnace  produces  about  50  tons  of  steel  in 


THE  IRON  FAMILY  347 

a  given  charge,  approximately  eight  hours  being  required  in 
the  process.  At  present  by  far  the  largest  amount  of  steel 
produced  in  the  United  States  is  made  by  this  process. 

Properties  of  steel.  Steel  contains  from  a  trace  up  to  2  per 
cent  of  carbon,  less  than  0.1  per  cent  of  silicon,  and  not  more 
than  traces  of  phosphorus  and  sulfur.  When  desired,  a 
product  containing  as  high  as  99.85  per  cent  iron  can  be 
produced  by  the  open-hearth  method.  Such  steel  is  very 
soft,  but  resists  rusting.  As  the  percentage  of  carbon  in- 
creases, the  steel  becomes  harder  and  less  ductile.  Steel  can 
be  rolled  into  sheets,  cast  in  molds,  and  forged  into  desired 
shapes. 

The  hardening  and  tempering  of  steel.  When  steel  con- 
taining from  0.5  to  1.5  per  cent  of  carbon  is  heated  to  a 
relatively  high  temperature  and  then  cooled  suddenly  by 
plunging  it  into  cold  water  or  oil,  it  becomes  very  hard  and 
brittle.  When  gradually  reheated  and  then  allowed  to  cool 
slowly,  this  hardened  steel  becomes  softer  and  less  brittle, 
and  this  process  is  known  as  tempering. 

By  properly  regulating  the  temperature  to  which  the  steel  is 
reheated  in  tempering,  it  is  possible  to  obtain  any  condition  of 
hardness  demanded  for  a  given  purpose,  as  for  making  springs 
or  cutting-tools.  Steel  assumes  different  color  tints  at  different 
temperatures,  and  by  these  the  experienced  workman  can  tell 
when  the  desired  temperature  has  been  reached.  Lake  gives 
the  following  temperatures  as  suited  to  the  tempering  of  the 
tools  specified : 

220° paper  cutters,  wood-engraving  tools 

240° knife  blades,  rock  drills 

200° hand-plane  cutters  and  cooper's  tools 

275° axes,  springs 

200° needles,  screw  drivers 

300°  wood  saws 


348  FIKST  COURSE  IK  CHEMISTRY 

Steel  alloys.  As  we  have  seen  (p.  339),  small  quantities 
of  carbon  greatly  modify  the  properties  of  iron,  and  equally 
marked  effects  may  be  produced  by  a  great  many  other 
elements.  Accordingly,  to  secure  a  steel  with  the  requisite 
properties,  suitable  percentages  of  these  elements  are  added 
to  the  steel  just  before  it  is  run  out  of  the  furnace.  The 
elements  most  frequently  added  are  manganese,  silicon, 
nickel,  chromium,  tungsten,  vanadium,  and  titanium,  and 
steel  containing  an  appreciable  percentage  of  any  of  these 
elements  is  called  a  steel  alloy.  The  element  is  added  in 
the  form  of  a  rich  alloy  of  iron,  such  as  ferrochromium  or 
ferromanganese. 

The  approximate  composition  and  uses  of  some  of  the 
principal  steel  alloys  is  as  follows: 

3.5%  nickel armor  plate 

3.5%  nickel  and  3.5%  chromium    .     .     .  armor  plate  and  projectiles 

4.0%  manganese     . burglar-proof  safes 

6.0%  chromium  and  from  8  to  24% 

tungsten high-speed  lathe  tools 

6.0%  chromium  and  10%  molybdenum  .  high-speed  lathe  tools 

0.1%  titanium car  rails  and  steel  castings 

0.1%  vanadium automobile  parts 

Steel  purifiers.  One  great  difficulty  in  securing  a  good 
steel  is  to  prevent  a  slight  oxidation  at  the  end  of  the  oper- 
ation, together  with  the  absorption  of  gases  which  cause 
blowholes  as  the  casting  solidifies.  These  difficulties  are 
avoided,  as  far  as  possible,  by  adding  to  the  steel,  at  the 
close  of  the  operation,  certain  elements  which  will  combine 
with  the  oxygen  and  the  absorbed  gases.  The  compounds 
formed  pass  into  the  slag,  and  almost  none  of  the  added 
element  remains  in  the  finished  product.  Aluminium  is 
used  to  a  large  extent  for  this  purpose  as  well  as  vanadium 
and  titanium.  Such  elements  are  called  purifiers. 


THE  IRON  FAMILY  349 

Compounds  of  iron.  Iron  differs  from  the  metals  so  far 
studied,  in  that  it  is  able  to  form  two  series  of  compounds. 
In  the  one  series  the  iron  is  bivalent  and  forms  compounds 
which  in  formulas  and  many  chemical  properties  are  similar 
to  the  corresponding-  zinc  compounds.  These  are  called  fer- 
rous compounds.  In  the  other  series  iron  acts  as  a  trivalent 
metal,  and  forms  salts  similar  to  those  of  aluminium.  These 
salts  are  known  as  ferric  compounds. 

Ferrous  salts.  These  salts  are  obtained  by  dissolving 
iron  in  the  appropriate  acid,  or,  when  insoluble,  by  pre- 
cipitation. The  crystallized  salts  are  usually  light-green 
in  color. 

Ferrous  sulfate  (FeS04).  Ferrous  sulfate  is  the  most 
familiar  ferrous  compound.  It  is  usually  obtained  in  the 
form  of  the  hydrate  FeSO4  •  7  H2O,  called  copperas,  or  green 
vitriol,  and  is  prepared  commercially  as  a  by-product  in 
the  steel-plate  mills.  Preparatory  to  galvanizing  or  tin- 
ning (p.  370),  steel  plates  are  cleaned  by  immersing  them 
in  dilute  sulfuric  acid,  and  in  the  process  some  of  the  iron 
dissolves.  The  liquors  are  concentrated,  and  the  green 
vitriol  separates  from  them.  The  salt  is  used  in  the  manu- 
facture of  ink  and  of  iron  alum,  and  as  a  reagent  to 
destroy  weeds. 

Ferrous  sulfide  (FeS).  Ferrous  sulfide  is  sometimes  found 
in  nature  as  a  golden-yellow  crystalline  mineral.  It  is  formed 
as  a  black  precipitate  when  a  soluble  sulfide  and  an  iron  salt 
are  brought  together  in  solution : 

FeS04  +  Na2S >•  FeS  +  Na2SO4 

It  can  also  be  made  as  a  heavy  dark-brown  solid  by  fusing 
together  the  requisite  quantities  of  sulfur  and  iron.  It  is 
used  in  the  laboratory  in  the  preparation  of  hydrogen  sul- 
fide (p.  145). 


350  FIKST  COURSE  IN  CHEMISTRY 

Iron  disulfide  (pyrite)  (FeS2).  This  substance  occurs  abun- 
dantly in  nature  in  the  form  of  brass-yellow  cubical  crystals 
and  in  compact  masses.  Sometimes  it  is  called  fool's  gold 
from  its  superficial  resemblance  to  the  precious  metal.  It  is 
used  in  very  large  quantities  as  a  source  of  sulfur  dioxide 
in  the  manufacture  of  sulfuric  acid,  since  it  burns  readily 
in  the  air,  forming  ferric  oxide  and  sulfur  dioxide  : 

4FeS  +11O  -  ^2FeO  +  8  SO 


23 


Ferrous  carbonate  (FeC03).  This  compound  occurs  in 
nature  as  siderite  and  is  a  valuable  ore.  Like  calcium 
carbonate,  it  dissolves  to  some  extent  in  water  containing 
carbon  dioxide,  and  waters  containing  it  are  called  cha- 
lybeate waters. 

Ferric  salts.  The  crystallized  ferric  salts  are  usually  yel- 
low or  violet  in  color.  Heated  with  water  in  the  absence  of 
free  acid,  they  hydrolyze  even  more  readily  than  the  salts 
of  aluminium.  The  most  familiar  ferric  salt  is  the  chloride. 

Ferric  chloride  (FeCl3).  This  salt  can  be  obtained  most 
conveniently  by  dissolving  iron  in  hydrochloric  acid  and 
then  passing  chlorine  into  the  solution  : 

Fe+2HCl  -  ^FeCl0  +  H2 

2FeCl  +  Cl 


The  crystallized  salt  has  the  formula  FeCl3  •  6  H0O. 

Ferric  hydroxide  (Fe(OH)3).  When  solutions  of  ferric 
salts  are  treated  with  ammonium  hydroxide,  ferric  hydrox- 
ide is  formed  as  a  rusty-red  precipitate  insoluble  in  water. 

Iron  rust  is  a  variable  mixture  of  hydrated  oxides  of  iron. 
When  a  film  of  rust  forms  on  iron  it  does  not  protect  the 
metal  from  the  further  action  of  water  as  does  the  rust  of 
aluminium  and  zinc,  because  iron  rust  is  porous  and  also 
tends  to  scale  off,  exposing  a  fresh  surface. 


THE  IKON  FAMILY  351 

Oxidation  of  ferrous  salts.  When  a  ferrous  salt  in  the 
presence  of  an  acid  is  oxidized  to  a  ferric  salt,  it  will  be 
noticed  that  the  valence  of  the  iron  is  increased  from  2  to  3. 
This  increase  in  valence  can  often  be  brought  about  without 
the  aid  of  oxygen,  as  is  shown  in  the  following  equation  : 

2  FeCl  4-  C1  -  >•  2  Fed 


, 


This  is  also  called  an  oxidation,  although  no  oxygen  takes 
part  in  the  reaction  (p.  20),  for  the  same  product  is  ob- 
tained as  by  the  other  method  : 

2  FeCl2  +  2  HC1  +  O  --  *  2  FeCl,  +  H2O 

In  general,  when  the  valence  of  the  metallic  ion  of  a  salt 
is  increased,  the  salt  is  said  to  be  oxidized,  whether  any 
oxygen  takes  part  in  the  reaction  or  not. 

Reduction  of  ferric  salts.  Ferric  salts  may  be  changed 
into  ferrous  salts  by  the  action  of  nascent  hydrogen  or 
other  reducing  agents,  as  shown  in  the  following  equations  : 

FeCl3  4-  [H]  --  >-  FeCl2  +  HC1 
2  FeCl  +  Zn  --  >-  2  FeCl  4-  ZnCl 

o  -- 

Although  no  oxygen  is  removed  (p.  31)  in  either  of  these 
reactions,  the  ferric  chloride  is  said  to  be  reduced;  and,  in 
general,  when  the  valence  of  the  metallic  ion  of  a  salt  is 
diminished,  the  salt  is  said  to  be  reduced. 

Potassium  ferrocyanide  (K4FeC6N6).  When  nitrogenous 
matter  such  as  horns  and  refuse  leather  is  heated  with 
potassium  carbonate  and  iron  borings,  and  the  mass  is  ex- 
tracted with  water,  there  crystallizes  from  the  solution  thus 
formed  a  beautiful  lemon-yellow  salt  of  the  composition 
K4FeC6N6  .  3  H2O.  This  is  called  potassium  ferrocyanide 
or  yellow  pru8#iat€  of  potash.  In  solution  it  gives  the  ions 
4K+  and  (FeC6N6)~~  "  but  no  ions  of  iron.  The  ion 


352  FIRST  COURSE  IN  CHEMISTRY 

(FeC6N6)~-  "  acts  as  a  radical,  and  many  different  ferro 
cyanides  can  be  obtained  by  precipitation.  For  example, 
a  ferric  salt  gives  an  intensely  blue  precipitate  of  ferric 
f  errocyanide : 

3  K/FeC6N6)  +  4  FeCl3 >-  Fe4(FeC6N6)8  +  12  KC1 

This  is  called  Prussian  blue.  It  is  used  as  a  paint  pigment 
and  for  bluing  laundry  water.  Unless  care  is  taken,  the 
alkali  of  the  soap  will  decompose  the  compound,  forming 
ferric  hydroxide  and  making  rust  stains. 

Potassium  ferricyanide  (K3FeC6N6).  By  treating  a  solu- 
tion of  potassium  ferrocyanide  with  chlorine  water  and 
evaporating  the  solution  to  crystallization,  garnet-red  crys- 
tals are  deposited  which  have  the  composition  K8FeC6N6 : 

2  K4FeC6N6  +  C12 >-  2  K8FeC6N6  +  2  KC1 

This  compound  is  called  potassium  ferricyanide,  or  red 
prussiate  of  potash. 

Blue  printing.  When  a  ferric  salt  and  potassium  ferricyanide 
are  brought  together  in  solution,  no  precipitate  forms,  though 
the  solution  acquires  a  yellowish  color.  On  exposure  to  the 
sunlight  the  ferric  salt  undergoes  a  partial  reduction  to  ferrous 
salt,  and  a  blue  precipitate  forms.  Advantage  is  taken  of  these 
facts  in  the  process  of  blueprinting.  A  sensitive  paper  is 
prepared  by  soaking  paper  in  a  solution  of  potassium  ferri- 
cyanide and  a  ferric  salt  (ferric  ammonium  citrate  is  generally 
used),  and  drying  it  in  a  dark  place.  When  a  black  drawing 
on  tracing  cloth  is  placed  upon  such  a  sensitive  paper  and  the 
two  are  exposed  to  the  sunlight,  the  sensitive  paper  (except 
where  it  is  protected  by  the  black  lines)  turns  a  brownish 
color.  It  is  then  thoroughly  washed  with  water,  to  remove  the 
soluble  salts,  during  which  process  the  portions  acted  upon  by 
the  light  turn  blue,  while  the  unaffected  portions  are  left  white. 
A  solution  of  sodium  hydroxide  can  be  used  as  an  ink  for  white 
lettering  on  a  blue  print,  since  this  base  decolorizes  Prussian  blue. 


THE  IRON  FAMILY  353 

Other  salts  of  iron.  The  following  compounds  of  iron 
have  industrial  uses : 

Ferric  sulfate  (Fe2(SO4)3)        a  white  solid 

Ferric  nitrate  (Fe(NO3)3  •  6  H2O)     .     .     .     violet  crystals 
Iron  alum  (NH4Fe(SO4)2  -  12  H2O)       .     .     violet  crystals 

Inks.  Most  of  the  common  black  inks  are  made  by  treating 
an  extract  of  nutgalls  with  ferrous  sulfate  and  adding  a  blue- 
black  dye.  The  nutgalls  are  rich  in  tannic  acid,  and  this,  with 
ferric  compounds  formed  by  the  oxidation  of  the  ferrous  sul- 
fate by  the  air,  gives  a  nearly  black  precipitate.  The  black  dye 
gives  a  temporary  color,  the  permanent  color  being  developed 
after  the  writing  has  been  exposed  to  the  air.  The  addition  of 
some  colloidal  material,  such  as  gum  arabic  or  dextrin,  together 
with  a  little  sulfuric  acid,  delays  the  precipitation  of  the  black 
substance  in  the  bottle.  A  preservative  is  usually  added,  to 
prevent  the  ink  from  molding. 

Removal  of  ink  and  other  stains.  Some  stains  may  be  removed 
by  methods  which  involve  no  applications  of  chemistry.  Thus, 
a  grease  spot  may  be  removed  by  placing  the  stain  over  some 
blotting  paper  and  washing  it  with  carbon  tetrachloride  (p.  207) 
or  benzine  (p.  205).  The  grease  is  dissolved  by  the  solvent,  and 
the  resulting  solution  is  absorbed  by  the  blotting  paper.  If  the 
grease  is  a  solid,  such  as  candle  grease  or  paraffin,  it  may  be 
removed  from  the  fabric  by  placing  the  stained  portion  between 
blotting  papers  and  pressing  it  with  a  hot  iron.  The  grease 
melts  and  is  absorbed  by  the  paper.  Turpentine  is  a  good 
solvent  for  paint  spots,  but  must  not  be  applied  to  silk.  Many 
substances  such  as  sirups  may  be  washed  out  with  water. 

In  many  cases  it  is  necessary  to  use  chemical  methods.  Thus, 
the  red  color  produced  by  many  acids  may  be  removed  by  wash- 
ing the  stained  portion  of  the  fabric  with  a  little  dilute  ammonia 
water.  Nitric  acid  acts  upon  the  cloth  as  well  as  upon  the  dye, 
so  that  the  original  color  cannot  be  restored.  Coffee  and  fruit 
stains  may  usually  be  removed  by  placing  the  stain  over  a  bowl 
and  pouring  boiling  water  upon  it.  If  not  removed  in  this  way, 
the  stain  may  be  washed  with  a  mild  bleaching  agent,  such  as 


354  FIRST  COURSE  IN  CHEMISTRY 

bleaching  powder,  to  which  some  water  and  a  few  drops  of  vine- 
gar have  been  added.  If  the  fabric  is  colored,  the  bleaching 
agent  may  act  upon  the  dye,  so  that  it  is  always  wise  in  such 
cases  first  to  try  the  effect  of  the  agent  upon  a  small  clipping 
of  the  fabric. 

Ink  stains  may  be  washed  out  with  water,  if  treated  at  once, 
as  in  the  case  of  fruit  stains.  When  the  ink  has  become  dry 
and  oxidized,  the  stain  may  be  removed  by  treating  with  lemon 
juice  or  with  a  dilute  solution  of  oxalic  acid.  By  this  treat- 
ment the  ferric  salts  in  the  ink  are  reduced  to  ferrous  salts 
which  can  then  be  washed  out  with  water.  Rust  stains  can  be 
removed  in  a  similar  way.  Some  indelible-ink  stains  may  be 
removed  by  soaking  the  fabric  in  a  solution  of  sodium  thiosul- 
fate.  Silk  is  so  sensitive  to  the  action  of  solvents  and  reagents  tlxit 
it  is  generally  impossible  to  remove  stains  from  it  without  injuring 
the  fabric. 

COBALT  AND  NICKEL 

Occurrence.  Cobalt  and  nickel  are  almost  always  found 
together  in  ores  which  also  contain  iron,  silver,  and  cop- 
per, in  combination  with  arsenic  arid  sulfur.  The  richest 
deposits  are  in  Ontario  and  New  Caledonia.  The  extraction 
of  these  metals  from  their  ores  and  their  separation  from 
each  other  is  too  complicated  a  process  to  be  described  here. 
Nickel  is  also  a  frequent  impurity  of  crude  copper,  and 
several  million  pounds  of  nickel  sulfate  are  annually  re- 
covered in  the  United  States  in  the  refining  of  copper  by 
electrolysis. 

Properties  and  uses.  Both  these  metals  are  silvery  in 
appearance  and  take  a  high  polish.  They  are  somewhat 
heavier  than  iron,  and  melt  at  a  lower  temperature.  Their 
chief  use  is  iu  making  alloys.  An  alloy  of  cobalt  and  chro- 
mium is  used  for  making  cutlery  and  lathe  tools.  Nickel 
coinage  consists  of  75  per  cent  copper  and  25  per  cent 
nickel.  German  silver  (p.  8(30)  also  contains  about  25  per 


THE  IKON  FAMILY 


355 


cent  nickel.  Nickel  is  extensively  used  as  a  plating  upon 
other  metals  (particularly  upon  brass),  to  prevent  tarnishing 
in  air,  and  cobalt  can  be  used  in  the  same  way. 

Electroplating  with  nickel.  Nickel  plating  is  accomplished 
by  an  electrolytic  process.  The  electrolyte  consists  of  a  solu- 
tion of  nickel  ammonium  sulfate,  a  salt  having  the  composition 
NiS04  •  (NH4)2S04  .  6  H20.  The  object  to  be  plated  is  sus- 
pended in  the  electrolyte  and  serves  as  the  cathode,  while  a 


FIG.  172.    Electroplating  with  nickel 

plate  of  nickel  is  used  as  the  anode.  When  the  current  is  pass- 
ing through  the  electrolyte,  the  nickel  is  deposited  upon  the  ob- 
ject to  be  plated,  and  an  equivalent  portion  of  nickel  dissolves 
from  the  anode,  the  composition  of  the  electrolyte  remaining 
unchanged.  Fig.  172  illustrates  the  process  carried  out  on  a 
large  scale,  the  objects  to  be  plated  being  suspended  from  the 
rods  A,  A. 

Cobalt  oxide  (CoO).  This  is  the  form  in  which  most  of  the 
cobalt  comes  into  the  market.  It  is  a  black  powder  used  in 
making  other  cobalt  compounds,  and  in  making  blue  glass  and 


356  FIRST  COUESE  IN  CHEMISTRY 

blue  decorations  on  china.    Sometimes  ground  blue  cobalt  glass, 
called  smalt,  is  used  instead  of  the  oxide,  and  as  a  pigment. 

Salts  of  cobalt  and  nickel.  Nearly  all  the  simple  salts  of 
cobalt  and  of  nickel  have  formulas  similar  to  those  of  ferrous 
salts.  The  most  familiar  are  the  following : 

Co(NO3)2  •  6  H2O       ....  a  cherry-red  deliquescent  salt 

CoCl2  •  6  H2O similar  in  appearance  to  the  nitrate 

CoS an  insoluble  black  precipitate 

KiSO4  •  7  H2O well-formed  green  crystals 

!Ni(ATO3)2  •  6  II2O deliquescent  green  crystals 

NiSO4  •  (NH4)2SO4  •  6  H2O      .  used  in  nickel  plating 

NiS an  insoluble  black  precipitate 

EXERCISES 

1.  Why  does  not  iron  occur  in  native  state?    What  does  its 
native  occurrence  in  meteors  indicate  ? 

2.  Why  is  the  furnace  in  which  cast  iron   is  made  called  a 
blast  furnace? 

3.  If  cast  iron  contained  no  carbon  or  silicon,  could  it  be  worked 
in  a  Bessemer  converter  ? 

4.  Why  is  the  air  heated  before  it  is  admitted  to  a  blast  furnace  ? 

5.  Write  equations  for  the  oxidation  of  ferrous  nitrate  to  ferric 
nitrate  in  the  presence  of  nitric  acid. 

6.  Write  equations  for  the  reduction  of  ferric  sulfate  by  nascent 
hydrogen. 

7.  Calculate  the  weight  of  iron  which  can  be  produced  from 
1  ton  of  hematite.    Ans.  1398.8  Ib. 

8.  What  is  the  meaning  of  the  term  magnetite!  '    . 

9.  What  are  some  of  the  advantages  of  plating  brass  with  nickel  ? 
10.    Calculate  the  percentage  of  iron  in  hematite,  in  magnetite, 

and  in  siderite.    .4n.s.  69.94;  71.92;  48.20. 

TOPICS  FOR  THEMES 

The  making  of  a  nail  (see  encyclopedia). 
The  making  of  a  needle  (see  encyclopedia). 
How  iron  castings  are  made.    (Visit  a  foundry.) 


CHAPTER  XXXVII 
COPPER,   MERCURY,   AND   SILVER 


FORMULAS  OF 

NAME 

SYMBOL 

ATOMIC 

DENSITY 

MELTING 

OXIDES 

ous 

ic 

Copper  . 

Cu 

63.57 

8.93 

1083° 

Cu2O 

CuO 

Mercury 

Ilg 

200.60 

13.56 

-  38.7° 

Hg,0 

HgO 

Silver    . 

Ag 

107.88 

10.50 

960.5° 

Ag20 

AgO 

The  family.  Although  mercury  is  not  in  the  same  family 
with  copper  and  silver,  the  three  elements  resemble  each  other 
so  closely  in  chemical  conduct  that  it  is  convenient  to  class 
them  together  for  study. 

COPPER 

Occurrence.  The  element  copper  has  been  used  for  vari- 
ous purposes  since  the  earliest  days  of  history.  It  is  often 
found  in  the  native  state,  large  masses  of  it  occurring  nearly 
pure  in  the  Lake  Superior  region  and,  to  a  smaller  extent, 
in  other  places.  The  most  valuable  ores  are  the  following : 


SULFUR  ORES 

Chalcopyrite 

Chalcocite      .... 

Bornite  Cu-FeS, 


CuFeSa 
Cu2S 


OXYGEN  ORES 
Cuprite   .     .     Cu2O 
Melaconite  .     CuO 
Malachite    .     CuCO3  •  Cu(OH)2 


Metallurgy  of  copper.  Ores  containing  little  or  no  sulfur 
are  easy  to  reduce.  They  are  first  crushed  and  the  earthy 
impurities  washed  away.  The  concentrated  ore  is  then 

357 


358 


FIRST  COURSE  IX  CHEMISTRY 


mixed  with  carbon  and  heated  in  a  furnace,  metallic  copper 
resulting  from  the  reduction  of  the  copper  oxide  by  the 
hot  carbon. 

Metallurgy  of  sulfur  ores.  Much  of  the  copper  of  commerce 
is  made  from  chalcopyrite  and  bornite,  and  these  ores  are  more 
difficult  to  work.  They  are  first  roasted  in  the  air,  by  which 


FIG.  173.    Mining  copper  ore  at  Butte,  Montana 

treatment  some  of  the  iron  is  converted  into  oxide.  Care  is 
taken,  however,  to  leave  enough  sulfur  to  combine  with  all  of 
the  copper  and  with  2^  part  of  the  iron. 

The  ore  so  treated  (or  coarse  ore  which  needs  no  treatment), 
together  with  a  flux  rich  in  silica,  is  charged  into  a  furnace 
called  a  matte  furnace.  In  this,  the  iron  oxide  combines  with 
the  silica  to  form  a  liquid  slag,  while  the  sulfides  of  iron  and 
copper  melt  into  a  heavier  liquid  called  matte. 

The  liquid  matte  is  then  tapped  off  into  a  converter  closely 
resembling  the  one  used  in  the  Bessemer  process  and  holding 


COPPER,  MERCURY,  AND  SILVER     359 

from  6  to  10  tons.  Some  silica  is  added,  and  a  carefully  regu- 
lated current  of  air  is  blown  in.  The  sulfur  acts  as  fuel,  burn- 
ing to  form  sulfur  dioxide ;  the  iron  sulfide  is  converted  into 
oxide,  which  then  combines  with  the  silica  to  form  slag;  the 
copper  sulfide  burns  to  form  sulfur  dioxide  and  copper.  When 
the  process  is  complete,  the  copper  is  poured  into  molds.  It  is 
called  blister  copper,  and  may  be  as  high  as  98  per  cent  pure. 
The  United  States  at  present  produces  about  1,500,000,000  Ib. 
of  copper  annually. 

Refining  of  copper.  Blister  copper  is  purified  by  electrol- 
ysis. A  large  plate  of  it,  "serving  as  an  anode,  is  suspended 
in  a  tank,  facing  a  thin  plate  of  pure  copper  which  is  the 
cathode.  The  tank  is  filled  with  a  solution  of  copper  sul- 
fate  and  sulfuric  acid  to  act  as  the  electrolyte.  A  current 
from  a  dynamo  passes  from  the  anode  to  the  cathode,  and 
the  copper,  dissolving  from  the  anode,  is  deposited  upon  the 
cathode  in  pure  form,  while  the  impurities  collect  on  the 
bottom  of  the  tank.  Electrolytic  copper  is  one  of  the  purest 
of  commercial  metals. 

Recovery  of  gold  and  silver.  Gold,  silver,  and  nickel  are  often 
present  in  small  quantities  in  copper  ores,  and  remain  in  the 
crude  copper.  In  electrolytic  refining  the  gold  and  silver  col- 
lect in  the  muddy  deposit  on  the  bottom  of  the  tank.  The  mud 
is  carefully  worked  over  from  time  to  time  and  the  precious 
metals  extracted  from  it.  A  surprising  amount  of  gold  and 
silver  is  obtained  in  this  way.  The  nickel  passes  into  solution 
and  is  recovered  from  the  electrolyte. 

Properties  of  copper.  Copper  is  a  rather  heavy  metal  of 
density  8.9,  and  has  a  characteristic  reddish  color.  It  is 
rather  soft,  and  is  very  malleable,  ductile,  and  flexible,  yet 
tough  and  strong;  it  melts  at  1083°  and  boils  at  2310°. 
As  a  conductor  of  heat  and  electrical  energy  it  is  second 
only  to  silver. 


360  FIRST  COURSE  IN  CHEMISTRY 

Since  it  is  below  hydrogen  in  the  displacement  series, 
hydrochloric  acid,  dilute  sulfuric  acid,  and  fused  alkalies 
are  almost  without  action  upon  it ;  nitric  acid  and  hot  con- 
centrated sulfuric  acid,  however,  readily  dissolve  it  (pp.  130, 
148).  In  moist  air  it  slowly  becomes  covered  with  a  film  of 
the  bright-red  oxide  Cu2O,  which  soon  changes  to  a  green 
basic  carbonate.  Heated  in  the  air  the  metal  is  easily  oxi- 
dized to  the  black  oxide  CuO. 

Uses.  Copper  is  extensively  used  for  electrical  purposes, 
for  roofs  and  cornices,  for  sheathing  the  bottoms  of  ships, 
and  for  making  alloys.  In  the  following  table  the  compo- 
sition of  some  of  these  alloys  is  indicated : 

Aluminium  bronze  .     .  90%-98%  copper,  2%-10%  aluminium 

Brass 63%-73%  copper,  27%~37%  zinc 

Bronze 70%-95%  copper,  l%-25%  zinc,  1%-IS%  tin 

German  silver      .     .     .  50%~60%  copper,  20%  zinc,  20%~30%  nickel 

Gun  metal 90%  copper,  10%  tin 

Gold  coin 10%  copper,  90%  gold 

Silver  coin 10%  copper,  90%  silver 

Nickel  coin      ....  75%  copper,  25%  nickel 

Electrotyping.  Books  are  often  printed  from  electrotype 
plates,  which  are  prepared  as  follows :  The  face  of  the  type 
is  covered  with  wax,  and  this  is  firmly  pressed  down  until  a 
clear  impression  is  obtained.  The  impressed  side  of  the  wax 
is  coated  with  graphite,  and  this  is  made  the  cathode  in  an  elec- 
trolytic cell  containing  a  copper  salt  in  solution.  The  copper  is 
deposited  as  a  thin  sheet  upon  the  letters  in  wax  and,  when 
detached,  is  a  perfect  copy  of  the  type,  the  under  part  of  the 
letters  being  hollow.  The  sheet  is  strengthened  by  pouring  on 
the  undersurface  a  suitable  amount  of  commercial  lead.  The 
sheet  so  strengthened  is  then  used  in  printing. 

Two  series  of  copper  compounds.  Copper,  like  iron,  forms 
two  series  of  compounds :  the  cuprous  compounds,  in  which 
it  is  univalent ;  and  the  cupric  compounds,  in  which  it  is 


COPPER,  MERCURY,  AND  SILVER  361 

bivalent.  The  cupric  salts  are  much  the  more  common  of 
the  two. 

Cuprous  compounds.  The  most  important  cuprous  com- 
pound is  the  oxide  Cu2O,  which  occurs  in  nature  as  ruby 
copper,  or  cuprite.  It  is  a  bright-red  substance  and  can  easily 
be  prepared  by  heating  copper  to  a  high  temperature  in  a 
limited  supply  of  air.  It  is  used  for  imparting  a  ruby  color 
to  glass.  By  treating  cuprous  oxide  with  different  acids  a 
number  of  cuprous  salts  can  be  made. 

Cupric  compounds.  Cupric  salts  are  easily  made  by  dis- 
solving cupric  oxide  in  acids,  or,  when  insoluble,  by  precipi- 
tation. In  crystallized  form  most  of  them  are  blue  or  green. 
Since  they  are  so  much  more  familiar  than  the  cuprous  salts, 
they  are  frequently  called  merely  copper  salts. 

Cupric  oxide  (CuO).  This  is  a  black  insoluble  substance 
obtained  by  heating  copper  in  excess  of  air,  or  by  igniting  the 
hydroxide  or  the  nitrate.  It  is  used  as  an  oxidizing  agent. 

Cupric  sulfate  (CuSOJ.  When  crystallized  from  water, 
copper  sulfate  forms  large  blue  crystals  of  the  hydrate 
CuSO4  •  5  H2O,  called  blue  vitriol,  or  bluestone.  The  salt  is 
a  by-product  in  silver  refining,  and  is  also  made  by  the 
oxidation  of  pyrite  containing  copper: 


The  salt  finds  extensive  use  in  electrotyping,  in  copper 
refining,  as  a  remedy  for  hoof  diseases  (particularly  in 
sheep),  and  in  the  manufacture  of  insecticides  (p.  255). 
Like  all  copper  salts,  it  is  poisonous,  especially  to  lower 
forms  of  life.  When  added,  even  in  very  minute  quantities, 
to  water  containing  green  pond  scum  (algae),  the  plant  is 
quickly  killed.  Mixed  with  milk  of  lime  (which  precipi- 
tates copper  hydroxide),  it  is  called  Bordeaux  mixture,  and 
is  used  as  a  spray  for  killing  molds  and  scale  on  fruit  trees. 


.362  FIRST  COUBSE  IN  CHEMISTRY 

Cupric  sulfide  (CuS).  In  the  form  of  a  black  insoluble 
precipitate  cupric  sulfide  (CuS)  is  easily  prepared  by  the 
action  of  hydrogen  sulfide  upon  a  solution  of  a  copper  salt: 

CuSO4  +  H2S  — >-  CuS  +  H2SO4 

MEHCUEY 

Occurrence.  Mercury  occurs  in  nature  chiefly  as  the  sul- 
fide HgS,  called  cinnabar.  The  mercury  mines  of  Spain 
have  long  been  famous,  and  California  is  the  next  largest 
producer. 

Metallurgy.  Mercury  is  a  volatile  metal  which  has  but 
little  affinity  for  oxygen,  and  this  makes  the  metallurgy  of 
mercury  very  simple.  The  crushed  ore,  mixed  with  a  small 
amount  of  carbon  to  reduce  any  oxide  or  sulf  ate  that  might 
be  formed,  is  roasted  in  a  current  of  air.  The  sulfur  burns 
to  sulfur  dioxide,  while  the  mercury  vaporizes  and  is  con- 
densed in  a  series  of  condensing  vessels.  The  metal  is  puri- 
fied by  distillation. 

Properties.  Mercury  is  a  heavy,  silvery  liquid,  with  a  den- 
sity of  13.56.  It  boils  at  357°  and  solidifies  at -38.7°.  It 
forms  alloys  (called  amalgams)  with  nearly  all  metals. 

Toward  acids  mercury  conducts  itself  very  much  like 
copper ;  it  is  easily  attacked  by  nitric  acid  and  by  hot 
concentrated  sulfuric  acid,  while  cold  sulfuric  acid  and 
hydrochloric  acid  have  no  effect  on  it. 

Uses.  Mercury  is  extensively  used  in  the  construction 
of  many  scientific  instruments,  such  as  the  thermometer 
and  the  barometer,  and  as  a  liquid  over  which  to  collect 
gases-  that  are  soluble  in  water.  The  readiness  with  which 
it  alloys  with  silver  and  gold  makes  it  very  useful  in  the 
extraction  of  these  elements.  All  salts  of  mercury  are  made 
directly  or  indirectly  from  the  purified  metal. 


COPPER,  MERCURY,  AND  SILVER     363 

Compounds  of  mercury.  Like  copper,  mercury  forms 
two  series  of  compounds  :  the  mereurom  compounds,  of 
which  mercurous  chloride  (HgCl)  is  an  example;  and 
the  mercuric  compounds,  represented  by  mercuric  chloride 

(Hgcy. 

Mercuric  oxide  (HgO).  Mercuric  oxide  is  usually  obtained 
as  a  brick-red  substance  by  carefully  heating  the  nitrate  : 


2  Hg  (N08)2  —  *2  HgO  +  4  NO2  +  O2 

It  can  also  be  obtained  in  a  yellow  form.  When  heated,  the 
oxide  darkens  until  it  becomes  almost  black  ;  at  a  higher 
temperature  it  decomposes  into  mercury  and  oxygen  (p.  4). 
Mercurous  chloride  (calomel)  (HgCl).  Being  insoluble, 
mercurous  chloride  is  precipitated  as  a  white  solid  when  a 
soluble  chloride  is  added  to  a  solution  of  mercurous  nitrate  : 

HgNO3  +  NaCl  --  >-  HgCl  +  NaNO3 

Commercially,  it  is  manufactured  by  heating  a  mixture  of 
mercuric  chloride  and  mercury.  It  is  a  common  medicine. 
Mercuric  chloride  (corrosive  sublimate)  (HgCl2).  This  sub- 
stance can  be  made  by  dissolving  mercuric  oxide  in  hydro- 
chloric acid.  On  a  commercial  scale  it  is  made  by  heating 
a  mixture  of  common  salt  and  mercuric  sulfate  : 

2  NaCl  +  HgSO4  --  >-  HgCl2  +  NaaSO4 

The  mercuric  chloride,  being  readily  volatile,  vaporizes,  and 
is  condensed  again  in  cool  vessels.  It  is  like  mercurous 
chloride  in  being  a  white  solid,  but  is  soluble  in  water.  It 
is  extremely  poisonous,  and  in  dilute  solutions  is  used  as 
an  antiseptic  in  dressing  wounds. 

Mercuric  sulfide  (HgS).  As  cinnabar,  this  substance  forms 
the  chief  native  compound  of  mercury,  and  occurs  in  red 
crystalline  masses.  By  passing  hydrogen  sulfide  into  a 


364:  FIKST  COUESE  IN  CHEMISTEY 

solution  of  a  mercuric  salt,  mercuric  sulfide  is  precipitated 
as  a  black  powder  insoluble  in  water  and  acids.  By  other 
means  it  can  be  prepared  as  a  brilliant  red  powder,  known 
as  vermilion,  which  is  used  as  a  pigment  in  fine  paints. 

SILVER 

Occurrence.  Silver  is  found  in  small  quantities  in  the 
uncombined  state ;  usually,  however,  it  occurs  in  combina- 
tion with  sulfur,  either  as  the  sulfide  Ag2S  or  as  a  constit- 
uent of  other  sulfides,  especially  those  of  lead  and  copper. 
It  is  also  found  alloyed  with  gold. 

In  this  country  silver  is  produced  almost  entirely  in  con- 
nection with  lead,  and  it  will  be  convenient  to  consider  the 
metallurgy  of  the  two  metals  together  in  the  next  chapter. 

The  refining  of  silver.  Crude  silver  obtained  by  any  proc- 
ess may  contain  a  number  of  metals,  especially  copper  and 
gold,  and  is  usually  refined  by  parting  with  sulfuric  acid. 
The  alloy  is  heated  with  concentrated  sulfuric  acid,  which 
dissolves  the  silver  and  copper,  but  not  the  gold.  In  the 
solution  of  silver  sulfate  so  obtained,  copper  plates  are  sus- 
pended upon  which  pure  silver  precipitates,  the  copper  going 
into  solution  as  the  sulfate,  as  shown  in  the  equation 

Ag2SO4  +  Cu  — >-  2  Ag  +  CuSO4 

The  solution  obtained  as  a  by-product  in  this  process  fur- 
nishes much  of  the  blue  vitriol  of  commerce.  Silver  is  also 
refined  by  electrolytic  methods  similar  to  those  used  in 
the  refining  of  copper. 

Properties  of  silver.  Silver  is  a  heavy,  rather  soft,  white 
metal,  very  ductile  and  malleable,  and  capable  of  taking  a 
high  polish.  It  surpasses  all  other  metals  as  a  conductor  of 
heat  and  electricity,  but  is  too  costly  to  find  extensive  use 


COPPER,  MERCURY,  AND  SILVER 


365 


for  such  purposes.  It  melts  at  a  little  lower  temperature 
than  copper.  It  alloys  readily  with  other  heavy  metals, 
and  when  it  is  to  be  used  for  coinage  or  for  tableware,  a 
small  amount  of  copper  —  from  8  per  cent  to  10  per  cent 
—  is  melted  with  it  to  give  it  hardness  (sterling  silver). 

It  is  not  acted  upon  by  water  or  air,  but  is  quickly 
tarnished  when  in  contact  with  sulfur  compounds  (eggs, 
mustard,  perspiration),  turning  quite  black  in  time.  Hy- 
drochloric acid  and  fused  alkalies  do  not  act  upon  it,  but 
nitric  acid  and  hot  concentrated  sulfuric  acid  dissolve  it 
with  ease.  When  a  solution  of  a  silver  salt  is  treated 

with  a  strong 
reducing  agent, 
metallic  silver 
is  precipitated. 
Under  proper 
conditions  this 
takes  the  form 
of  a  brilliant 

mirror  deposited  on  the  sides  of  the  glass  vessel.    Mirrors 
are  usually  made  in  this  way. 

Electroplating  with  silver.  Since  silver  is  not  acted  upon  by 
water  or  air,  and  has  a  pleasing  appearance,  it  is  used  to  coat 
various  articles  made  of  cheaper  metals.  Such  articles  are  said 
to  be  silver  plated,  and  the  process  by  which  this  is  done  is  very 
similar  to  electroplating  with  nickel  (p.  355).  The  object  to 
be  plated  (as,  for  example,  a  spoon)  is  attached  to  a  wire  and 
dipped  into  a  solution  of  a  suitable  silver  salt.  Electrical  con- 
nection is  made  in  such  a  way  that  the  article  to  be  plated  is 
the  cathode  (Fig.  174),  while  the  anode  A  is  made  up  of  one 
or  more  plates  of  silver. 

Compounds  of  silver.  Silver  forms  only  one  series  of  salts, 
which  corresponds  to  the  mercurous  and  the  cuprous  series. 


FIG.  174.    The  process  of  silver  plating 


366  FIRST  COURSE  IN  CHEMISTRY 

Silver  nitrate  (lunar  caustic)  (AgN03).  This  salt  is  easily 
prepared  by  dissolving  silver  in  nitric  acid,  and  evaporating 
the  resulting  solution.  It  crystallizes  in  flat  colorless  plates, 
and  when  heated  carefully  can  be  melted  without  decom- 
position. When  cast  into  sticks  it  is  called  lunar  caustic, 
for  it  has  a  very  corrosive  action  on  flesh,  and  is  sometimes 
used  in  surgery  to  burn  away  abnormal  growths. 

The  alchemists  designated  the  metals  by  the  names  of  the 
heavenly  bodies.  The  moon  (linut')  was  the  symbol  for  silver; 
hence  the  name  hnmr  <-<nixti<-. 

Silver  sulfide  (Ag2S).  This  occurs  in  nature  and  consti- 
tutes one  of  the  principal  ores  of  silver.  It  can  be  obtained 
as  a  black  solid  by  heating  silver  and  sulfur  together  or  by 
passing  hydrogen  sulfide  into  a  solution  of  silver  nitrate. 

Compounds  of  silver  with  the  halogens.  The  chloride, 
bromide,  and  the  iodide  of  silver  are  insoluble  in  water  and 
in  acids,  and  therefore  are  precipitated  by  bringing  together 
a  soluble  halogen  salt  with  silver  nitrate : 

AgX03  +  KC1 >•  AgCl  +  KN03 

They  are  remarkable  for  the  fact  that  they  are  very  sensi- 
tive to  the  action  of  light,  undergoing  a  change  of  color 
and  chemical  composition  when  exposed  to  sunlight,  espe- 
cially if  in  contact  with  organic  matter,  such  as  gelatin.  It 
is  upon  this  property  of  the  silver  halides  that  the  art  of 
photography  is  based. 

Photography.  From  a  chemical  standpoint  the  processes  of 
photography  may  be  described  under  two  heads :  (1)  the  prep- 
aration of  the  negative  ;  (2)  the  preparation  of  the  print. 

1.  Preparation  of  the  negative.  The  plate  used  in  the  prepara- 
tion of  the  negative  is  made  by  spreading  a  thin  layer  of  gela- 
tin, in  which  colloidal  silver  bromide  is  suspended  (silver 
iodide  is  sometimes  added  also),  over  a  glass  plate  or  celluloid 


COPPER,  MEKCURY,  AND  SILVER 


367 


FIG.  175.   The  negative  plate 


film  and  allowing  it  to  dry.  When  the  plate  so  prepared  is 
placed  in  a  camera  and  the  image  of  some  object  is  focused 
upon  it,  the  silver  salt  undergoes  a  change  which  is  propor- 
tional at  each  point  to  the 
intensity  of  the  light  fall- 
ing upon  it.  In  this  way 
an  image  of  the  object 
photographed  is  produced 
upon  the  plate.  This  im- 
age, however,  is  invisible 
and  is  therefore  called 
latent.  It  can  be  made 
visible  by  the  process  of 
developing. 

To  develop  the  image  the  exposed  plate  is  immersed  in  a 
solution  of  some  reducing  agent  called  the  developer.  While 
the  developer  will  in  tim,e  reduce  all  the  silver  salt,  it  acts 
much  more  rapidly  upon  that  which  has  been  exposed  to  the 
light.  The  action  is  therefore  continued  only  long  enough  to 
bring  out  the  image.  The  reduced  silver  is  deposited  in  the 
form  of  a  black  film  which  adheres  closely  to  the  plate. 

The  unaffected  silver  salt  is  now  removed  from  the  plate  by 
immersing  it  in  a  solution  of  sodium  thiosulfate  (hypo).  The 
plate  is  then  washed  with 
water  and  dried.  The  plate 
so  prepared  is  called  the 
negative  (Fig.  175),  be- 
cause it  is  a  picture  of 
the  object  photographed, 
with  the  lights  and  posi- 
tions exactly  reversed. 

2.    Preparation    of    the 
print.    The  print  is  made 
on    paper    which   is   pre- 
pared much  in  the  same  way  as  the  negative  plate.    The  nega- 
tive is  placed  upon  this  paper  and  exposed  to  the  light  in  such 
a  way  that  the  light  must  pass  through  the  negative  before 
striking  the  paper.   If  the  paper  is  coated  with  silver  chloride, 


FIG.  170.    The  positive  print 


368  FIRST  COUESE  IN  CHEMISTRY 

a  visible  image  is  produced,  in  which  case  a  developer  is  not 
needed.  Proofs  are  made  in  this  way.  In  order  to  make  them 
permanent,  the  unchanged  silver  chloride  must  be  dissolved  off 
with  sodium  thiosulfate.  The  print  is  then  toned  by  dipping  it 
into  a  solution  of  gold  or  platinum  salts,  in  which  process  the 
silver  on  the  print  passes  into  solution,  while  the  gold  or  plati- 
num takes  its  place.  These  metals  give  a  characteristic  color 
or  tone  to  the  print,  the  gold  making  it  reddish  brown,  while 
the  platinum  gives  it  a  steel-gray  tone.  Since  the  darkest  places 
on  the  negative  cut  off  the  most  light,  it  is  evident  that  the 
lights  of  the  print  (Fig.  176)  will  be  the  reverse  of  those  of 
the  negative,  and  will  therefore  correspond  to  those  of  the 
object  photographed. 

EXERCISES 

1.  Why  has  copper  or  bronze  been  used  for  so  long  a  time? 

2.  Why  do  we  have  so  many  relics  from  the  bronze  age  and  so 
few  from  the  iron  age? 

3.  Why  is  a  solution  of  copper  sulfate  acid  to  litmus  paper? 

4.  How  would  you  account  for  the  fact  that  so  many  different 
salts  of  copper  have  the  same  blue  color  in  dilute  solutions  ? 

5.  How  could  you  distinguish  between  mercurous  chloride  and 
mercuric  chloride? 

6.  Crude  silver  usually  contains  iron  and  lead.    What  would 
become  of  these  in  refining  silver  by  parting  with  sulfuric  acid  ? 

7.  Mercuric  nitrate  and  silver  nitrate  are  both  white  soluble 
solids.    How  could  you  distinguish  between  them  ? 

8.  How  do  you  account  for  the  fact  that  a  silver  spoon  grad- 
ually darkens  when  in  contact  with  eggs  ? 

9.  Why  are  all  three  of  these  metals  found  to   some  extent 
native  in  nature  ? 

10.    What  properties  make  mercury  useful  in  thermometers  and 
in  barometers  ? 

TOPICS  FOR  THEMES 

The  history  of  a  Kodak  picture  (Lassar-Cohn,  Chemistry  in  Daily 
Life). 

How  mirrors  are  made  (Lassar-Cohn,  Chemistry  in  Daily  Life). 
How  thermometers  are  made  (see  encyclopedia ). 


CHAPTER  XXXVIII 
TIN  AND  LEAD 


NAME 

SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POINT 

COMMON  OXIDES 

Tin  .      . 
Lead 

Sn 
Pb 

119.0 
207.1 

7.3 
11.37 

232° 

327° 

SnO                  SnO2 
PbO   Pb3O4  PbO2 

TIN 

Occurrence.  Tin  is  found  in  nature  chiefly  as  the  oxide 
SnO2,  called  cassiterite,  or  tinstone.  The  most  famous  mines 
are  in  Cornwall,  in  the  East  Indies,  and  in  Bolivia.  Very 
little  is  produced  in  the  United  States. 

Metallurgy.  The  metallurgy  of  tin  is  very  simple.  The 
ore,  separated  as  far  as  possible  from  earthy  materials,  is 
mixed  with  carbon  and  heated  in  a  furnace,  the  reduction 
taking  place  readily.  The  equation  is 

SnO2  +  C  — >•  Sn  +  CO2 

The  metal  is  often  purified  by  carefully  heating  it  until 
it  is  partly  melted ;  the  pure  tin  melts  first  and  can  be 
drained  away  from  the  impurities. 

Properties.  Pure  tin,  called  block  tin,  is  a  soft  white 
metal  with  a  silver-like  appearance  and  luster;  it  melts 
readily  (232°)  and  is  somewhat  lighter  than  copper,  having 
a  density  of  7.3.  It  is  malleable  and  can  be  rolled  out  into 
very  thin  sheets,  forming  tin  foil ;  most  tin  foil,  however, 
contains  a  marked  percentage  of  lead. 

369 


370  FIKST  COUKSE  IN  CHEMISTRY 

Under  ordinary  conditions  tin  is  unchanged  by  air  or 
moisture,  but  at  a  high  temperature  it  burns,  forming  the 
oxide  SnO2.  Dilute  acids  have  little  effect  upon  it,  but 
concentrated  acids  attack  it  readily.  Concentrated  hydro- 
chloric acid  changes  it  into  the  chloride  : 


Sn  +  2  HC1  --  >-  SnCl  +  H 


With    sulfuric    acid,   tin    sulfate   and   sulfur    dioxide   are 
formed  : 

Sn  +  2  H2S04  -  >-  SnS04  +  SO2  +  2  H2O 

Concentrated  nitric  acid  oxidizes  it,  forming  a  white  insolu- 
ble compound,  H2SnOg,  called  metastannic  acid. 

Uses  of  tin.  A  great  deal  of  tin  is  used  in  the  making 
of  tin  plate.  The  process  consists  in  dipping  thin  sheets  of 
iron  into  the  melted  tin  and  is  quite  similar  to  that  of 
galvanizing  iron  (p.  319).  Owing  to  its  resistance  to  the 
action  of  air  and  weak  acids,  tin  plate  is  used  in  many  ways, 
such  as  in  roofing,  and  in  the  manufacture  of  tin  cans, 
cooking  vessels,  and  similar  articles.  Small  pipes  of  block 
tin  are  used  instead  of  lead  for  conveying  pure  water  or 
liquids  containing  dilute  acids,  such  as  soda  water.  Many 
useful  alloys  contain  tin  (p.  360).  Pewter  and  soft  solder 
are  alloys  of  tin  and  lead. 

Soldering  and  brazing.  The  use  of  solder  in  joining  two  metal 
surfaces  depends  upon  (1)  the  low  melting  point  of  the  solder, 
and  (2)  the  fact  that  it  flows  over  clean  metal  surfaces  and 
sticks  to  them  on  cooling.  To  secure  clean  surfaces  free  from 
oxide,  a  suitable  flux  must  be  used  which  will  either  flixso/re 
the  oxide  as  fast  as  it  forms,  or  will  reduce  it  again  to  metal. 
The  usual  fluxes  are  zinc  chloride,  ammonium  chloride,  rosin, 
and  stearin.  In  brazing  or  7/^/v/  sotilci'ing  the  process  is  essen- 
tially the  same,  except  that  a  low-melting  brass  is  used  instead 
of  solder,  and  borax  is  used  as  a  flux. 


TIN  AND  LEAD  371 

Compounds  of  tin.  Tin  forms  two  series  of  metallic  com- 
pounds:  the  xtunHtmta,  in  which  the  tin  is  bivalent,  as  is 
illustrated  in  the  compounds  SnO,  SnS,  SnCl0;  and  the 
stannic,  in  which  it  is  tetravalent,  as  shown  in  the  com- 
pounds SnO.,,  SnS.,. 

Chlorides  of  tin.  Stannous  chloride  is  prepared  by  dis- 
solving tin  in  concentrated  hydrochloric  acid  and  evaporat- 
ing the  solution  to  crystallization.  The  crystals  which  are 
obtained  have  the  composition  SnCl,,  •  2  H2O,  and  are  known 
as  tin  crystals.  By  treating  a  solution  of  stannous  chloride 
with  aqua  regia,  stannic  chloride  is  formed: 

SnCl2  +  2  Cl >-  SnCl4 

The  salt  which  crystallizes  from  such  a  solution  has  the 
composition  SnCl4  •  5  H.,O,  and  is  known  commercially  as 
oxymuriate  of  tin.  If  metallic  tin  is  heated  in  a  current  of 
dry  chlorine,  anhydrous  stannic  chloride  (SnCl4)  is  obtained 
as  a  heavy,  colorless  liquid  which  fumes  strongly  on  ex- 
posure to  air.  A  great  deal  of  tin  in  the  form  of  stannic  chlo- 
ride is  recovered  from  scrap  tin  by  the  action  of  chlorine. 

The  crystallized  chlorides  of  tin  are  much  used  as  mor- 
dants in  dyeing  processes,  and  in  calico  printing. 

LEAD 

Occurrence.  Lead  is  found  in  nature  chiefly  as  the  sul- 
fide  PbS,  called  galenite  (Fig.  178).  In  the  United  States 
this  is  mined  principally  in  Missouri  (Fig.  177)  and  Idaho. 

Metallurgy  of  lead.  Almost  all  the  lead  of  commerce 
is  made  from  galenite,  which  usually  contains  some  silver. 
To  obtain  this  silver  most  economically  it  is  customary  to 
combine  richer  silver  ores  with  lead  ores  and  work  the 
two  together. 


372 


FIEST  COURSE  IN  CHEMISTRY 


Reduction  of  silver-bearing  lead.  The  sulfide  ores  are 
first  roasted  until  a  part  of  the  sulfide  has  been  changed 
into  the  oxide  and  the  sulfate.  The  air  is  then  shut  off 
and  the  heating  continued,  which  brings  about  the  reactions 
indicated  in  the  following  equations : 

2  PbO  +  PbS >•  3  Pb  +  SO2 

PbS04  +  PbS  — >•  2  Pb  +  2  S02 

By  reactions  which  are  similar  to  the  above,  the  ores  bear- 
ing silver  are  reduced,  the   silver  alloying  with  the  lead. 


•FiG.  177.   Mining  galenite  in  the  Joplin  region  in  Missouri 

The  softening  of  lead.  The  lead  obtained  in  this  way  is 
called  hard  lead,  and  in  addition  to  silver  contains  smaller 
quantities  of  other  elements,  especially  of  copper,  arsenic, 
antimony,  gold,  and  bismuth.  The  lead  is  softened  by  melt- 
ing it  in  an  open  furnace  with  free  access  of  air.  This  con- 
verts most  of  the  impurities  (as  well  as  some  lead)  into 
oxides  which  float  upon  the  melted  lead  and  can  easily  be 
removed.  The  partially  purified  lead  is  called  soft  lead. 


TIN  AND  LEAD  373 

Desilverizing  of  lead  (the  Parkes  process).  The  lead  is  melted 
in  large  kettles  holding  as  much  as  30  tons,  and  about  1  per  cent 
of  zinc  is  stirred  in.  These  two  metals  do  not  mix  to  any  great 
extent,  and  gold,  silver,  and  copper  are  much  more  soluble  in 
zinc  than  in  lead.  Consequently  when  the  stirring  ceases,  the 
zinc  rises  to  the  surface  of  the  lead,  carrying  with  it  the  other 
metals.  The  zinc  is  then  skimmed  off  and  distilled  (by  which 
process  the  zinc  is  recovered  to  be  used  again),  and  the  residue 
of  silver  and  gold  is  melted  down  and  cast  into  ingots  called 
dore  bars.  These  are  refined  as  explained  under  silver  (p.  364). 
An  electrolytic  method  (Belts  process)  is  now  being  used, 
similar  to  the  one  employed  with  copper,  but  with  many 
special  details. 

Properties.  Lead  is  a  heavy  metal  which  has  a  brilliant 
silvery  luster  on  a  freshly  cut  surface,  but  which  soon  tar- 
nishes by  oxidation  to  a  dull  blue-gray  color.  It  is  soft, 
easily  fused  (melting  at  327°),  and  malleable,  but  has  little 
toughness  or  strength. 

It  is  not  acted  upon  to  any  great  extent,  under  ordinary 
conditions,  by  the  oxygen  of  the  air,  but  at  a  high  tempera- 
ture is  changed  into  the  oxide.  With  the  exception  of 
hydrochloric  and  sulfuric  acids  (which  form  insoluble  com- 
pounds), most  acids,  even  very  weak  ones,  act  upon  it, 
forming  soluble  lead  salts.  Hot  concentrated  hydrochloric 
and  sulfuric  acids  also  attack  it  to  a  slight  extent. 

Uses.  Lead  finds  many  important  applications  in  the  in- 
dustries, chiefly  in  the  manufacture  of  storage  batteries,  in 
linings  for  sulfuric  acid  plants,  in  alloys  of  various  kinds 
(such  as  shot,  antifriction  metals,  type  metal,  and  pewter), 
and  in  water  pipes  for  plumbing.  Since  lead  dissolves  to 
some  extent  in  pure  water,  it  should  not  be  used  for 
pipes  that  are  to  carry  rain  water.  About  one  third  of  the 
annual  production  of  lead  is  used  in  making  paint,  and  is 
permanently  lost. 


374  FIRST  COURSE  IN  CHEMISTRY 

Compounds  of  lead.  In  nearly  all  its  compounds  lead  is 
bivalent,  but  in  a  few  of  its  compounds  it  has  a  valence  of 
four.  All  of  its  compounds  are  poisonous. 

Lead  oxides.  Lead  forms  a  number  of  oxides,  the  most 
important  of  which  are  the  following : 

1.  Litharge  (PbO).  This  oxide  forms  when  lead  is  oxi- 
dized at  a  rather  low  temperature,  and  is  obtained  as  a 


FIG.  178.    A  crystal  of  galenite  embedded  in  calcite 

by-product  in  silver  refining.  It  apparently  exists  in  a 
number  of  different  forms  ranging  from  yellow  and  light 
brown  to  red.  It  has  a  number  of  commercial  uses. 

2.  Red  lead,  or  minium  (P6304).  Minium  is  prepared  by 
heating  lead  (or  litharge)  to  a  high  temperature  in  contact 
with  a  current  of  air.  It  is  a  heavy  powder  of  a  beautiful 
red  color  and  is  much  used  as  a  pigment  for  painting 
structural  iron.  Mixed  with  linseed  oil  it  forms  a  cement 
used  in  joining  gas  pipes. 


TIN  AND  LEAD 


375 


3.  Lead  peroxide  (Pb02).  This  is  left  as  a  residue  when 
minium  is  heated  with  nitric  acid.  It  is  a  brown  powder 
which  easily  gives  up  a  part  of  its  oxygen  and  is  a  good 
oxidizing  agent. 

Lead  sulfide  (PbS).  In  nature  this  compound  occurs  in  a 
highly  crystalline  form  called  galenite  (Fig.  178),  the  crystals 
having  much  the  same  color  and  luster  as  pure  lead.  It  is 
readily  prepared  in  the  laboratory  as  a  black  precipitate, 
by  the  action  of  hydrogen  sulfide  upon  soluble  lead  salts  : 


HS 


PbS  +  2  HN0 


Pb(NOa)2 

It  is  insoluble  both  in  water  and  in  dilute  acids. 

Lead  carbonate.    While   the  normal  carbonate   of  lead, 
PbCOQ,  is  found   to  some  extent  in  nature  and  can  be 

o 

prepared  in  the  laboratory,  basic  car- 
bonates of  varying  composition  are 
much  more  easy  to  obtain.  One  of 
the  simplest  of  these  has  the  com- 
position (PbCO3)2  •  Pb(OH)2,  and  is 
called  white  lead.  This  is  prepared 
on  a  large  scale  as  a  white  pigment 
and  as  a  body  for  paints  which  are 
to  be  colored  with  other  substances. 

Manufacture  of  white  lead.  \Yhite  lead 
can  be  prepared  by  a  number  of  proc-     FIG    1?9>    A  crock  CQn_ 
esses,  but  no  other  seems  to  produce     taining   lead   plates   for 
a    product    of    as    desirable    physical          making  white  lead 
properties   as  the   old  Dutch  process, 

which  has  been  used  for  centuries,  though  with  many  improve- 
ments. In  this  process  the  lead  is  cast  into  perforated  plates 
called  buckles,  which  are  placed  loosely  upon  each  other  in  a 
crock  of  the  shape  shown  in  Fig.  179,  the  ledge  B  formed  by 
the  constriction  of  the  crock  supporting  the  plates.  Under 
them  in  A  is  poured  a  suitable  quantity  of  dilute  acetic  acid, 


376  FIRST  COUESE  IN  CHEMISTEY 

and  the  crocks  so  charged  are  placed  in  banks  arid  covered  with 
stable  manure  or  spent  tanbark.  The  heat  of  fermentation  in  the 
•latter  warms  the  acid,  the  fumes  of  which  attack  the  lead, 
forming  lead  acetate.  The  carbon  dioxide  from  the  fermentation 
enters  into  reaction  with  the  acetate  and  produces  the  basic 
carbonate,  regenerating  acetic  acid,  which  acts  again  upon  the 
lead.  The  process  continues  until  the  buckles  are  almost  com- 
pletely converted  into  the  desired  compound.  Fig.  180  shows  a 

buckle  before  and  after 
the  corrosion. 

Paints.  A  paint  con- 
sists of  three  essential 
ingredients:  the  vehi- 
cle, the  body,  and  the 
pigment. 

FIG.  180.    Lead  buckle  before  and  after  L   The  vehicle>  or  li*uid 

exposure  to  acetic  acid  and  carbon  dioxide      medium.     Ihis    must   be 

an   oil   which  will   dry 

rapidly  and  harden  in  drying  to  a  more  or  less  flexible,  horn- 
like body.  These  changes  in  the  oil  are  due  to  oxidation  by 
the  air.  A  number  of  different  oils  will  serve  this  purpose, 
but  linseed  oil  has  long  been  used  as  the  standard  drying  oil, 
since  it  can  be  produced  in  quantity  and  at  moderate  cost.  It  is 
customary  to  add  to  it  a  dryer,  made  by  boiling  some  of  the 
oil  with  oxides  of  manganese,  lead,  or  cobalt.  The  oxides  enter 
into  combination  with  the  oil  and  assist  catalytically  in  its 
oxidation. 

2.  The  body.  The  body  of  the  paint  must  be  some  solid  mate- 
rial, suspended  in  the  oil,  which  will  give  a  smooth  and  waxy 
surface  as  the  paint  dries,  and  will  have  good  covering  power. 
While  white  lead  meets  these  requirements,  it  is  moderately 
expensive  and  it  also  blackens  when  exposed  to  hydrogen  sul- 
fide,  which  is  likely  to  be  present  in  the  air  in  cities.  Other 
bodies  are  now  frequently  combined  with  the  lead,  or  replace  it 
altogether,  among  them  being  zinc  oxide,  China  clay  (or  kaolin), 
barium  sulfate,  and  a  product  called  lithopoiie  (p.  320).  For  some 


TIN  AND  LEAD 


37T 


purposes  these  materials  are  a  real  advantage,  and  they  are 
not  to  be  regarded  as  adulterants  unless  sold  as  white  lead. 
3.  The  pigment,  or  coloring  matter.  In  the  case  of  white  paints 
the  body  serves  also  as  the  coloring  matter.  For  other  colors 
a  specific  pigment  must 
be  added.  In  most  cases 
these  pigments  are  me- 
tallic oxides  or  salts,  and 
are  frequently  natural 
products.  Sometimes 
they  are  prepared  by 
precipitating  an  amor- 
phous body  (usually  a 
colloid)  in  the  presence 
of  an  organic  dye,  the 
dye  being  absorbed  by 
the  precipitate  and  giv- 
ing it  a  color.  Such  pig- 
ments can  be  prepared 
in  endless  variety  of  col- 
ors, and  are  called  lakes. 
They  are  usually  not  so 
permanent  as  mineral 


FIG.  181.   Apparatus  used  in  grinding  and 
mixing  paints 


pigments. 

Fig.  181   represents 
the  method  of  manufac- 
ture of  paint.  The  body,  together  with  a  little  oil,  enters  at  A  and 
is  ground  in  succession  in  B,  C,  D,  and  E,  during  which  process 
the  requisite  amounts  of  oil,  dryer,  and  pigment  are  added. 


OTHER  IMPORTANT  COMPOUNDS  OF  LEAD 

Lead  nitrate  (Pb(NO3)2)  :  white  soluble  crystals 

Lead  chloride  (PbCl2)  :  white  needles,  very  sparingly  soluble 

Lead  sulfate  (PbSO4)  :  an  insoluble  white  crystalline  powder 

Lead  acetate  (Pb(C2H8O2)2  •  3  H2O)  :    a   soluble  white   salt  called 

sugar  of  lead 
Lead  chromate  (PbCrO4)  :  used  as  a  pigment  in  paint  (chrome  yellow*) 


3T8  FIRST  COURSE  IN  CHEMISTRY 

EXERCISES 

1.  How  could  you  detect  lead  if  it  were  present  in  tin  foil? 

2.  Stannous   chloride   reduces  gold   chloride   (AuCl8)  to   gold. 
Give  equation. 

3.  What  are  the  products  of  hydrolysis  when  stannic  chloride  is 
used  as  a  mordant? 

4.  How  could  you  detect  arsenic  or  copper  in  lead? 

5.  In  obtaining  stannic  chloride  from  scrap  tin  plate  the  product 
is  contaminated  with  ferric  chloride.     Can  you  suggest  a  way  to 
purify  it? 

6.  What  sulfates  other  than  lead  are  insoluble? 

7.  Could  lead  nitrate  be  used  in  place  of  barium  chloride  in 
testing  for  sulfates? 

8.  The  purity  of  white  lead  is  usually  determined  by  observing 
the  weight  of  carbon  dioxide  given  off  when  it  is  treated  with  an 
acid.   On  the  supposition  that  it  has  the  formula  (PbCO3)2  •  Pb(OII)0, 
how  nearly  pure  was  the  sample  if  1  g.  gave  102.1  mg.  of  carbon 
dioxide?   Ans.  90%  pure. 

9.  Silicon  belongs  in  the  same  family  with  tin  and  lead.    In 
what  respects  are  these  elements  similar  ? 

10.  What  reaction  would  you  expect  to  take  place  when  lead 
peroxide  is  treated  with  hydrochloric  acid  ? 

11.  What  weight  of  tin  could  be  obtained  by  the  reduction  of 
1  ton  of  cassiterite  ?    Ans.  1576  Ib. 

12.  White  lead  is  often  adulterated  with  barite.  Suggest  a  method 
for  detecting  it,  if  present,  in  a  given  specimen  of  white  lead. 

TOPICS  FOR  TlIKMKS 

The  lead  mines  of  Missouri   (Geological   Reports,   Washington 
Bureau  of  Mines).  • 

Tin  mining  in  Cornwall  (see  encyclopedia). 
How  shot  is  made  (see  encyclopedia). 


CHAPTER  XXXIX 
URANIUM  AND  RADIUM 

Uranium.  Uranium  is  a  rare  element  whose  compounds 
were  first  isolated  from  a  mineral  called  pitchblende  or 
uraninite,  which  is  essentially  an  oxide  of  the  formula 
U3Og.  Carnotite,  a  mineral  discovered  more  recently  (1899), 
contains  both  uranium  and  vanadium,  and  is  found  chiefly 
in  Colorado  and  Utah.  The  carnotite  ores  are  by  far  the 
most  abundant  source  of  uranium,  the  next  in  importance 
being  the  pitchblende  deposits  in  Austria,  owned  by  the 
government.  The  American  production  of  uranium  ores  in 
1912  amounted  to  over  twice  as  much  as  that  of  all  other 
countries  combined. 

Compounds  and  uses.  The  most  familiar  compounds  of 
uranium  are  the  black  oxide,  U3Og,  and  the  yellow  uranyl 
nitrate,  UO0(NO8)2.  Most  of  the  compounds  of  uranium 
are  yellow  or  red.  Their  chief  chemical  use  is  in  making 
greenish-yellow  fluorescent  glass,  in  the  decorating  of  china 
with  various  shades  of  yellow,  orange,  and  black,  and  in 
the  making  of  orange-colored  pigments.  Uranium  steel  alloy 
has  useful  properties. 

Radioactivity  of  uranium.  In  1896  the  French  physicist 
Becquerel  discovered  that  uranium  and  all  its  compounds 
possess  a  property  which  has  been  named  radioactivity. 
This  radioactivity  manifests  itself  in  the  following  ways: 
(1)  A  photographic  plate  wrapped  in  black  paper  and 
placed  near  a  compound  of  uranium  is  affected  as  though 

379 


380 


FIRST  COURSE  IN  CHEMISTRY 


exposed  to  light.  A  metallic  object  placed  on  the  plate 
screens  the  plate  from  this  action,  and  leaves  its  outline 
on  the  plate  when  it  is  developed,  forming  a  radiograph 


FIG.  182.    A  radiograph  of  some  metal  objects 

(Fig.  182).  (2)  A  charged  electroscope  is  rapidly  dis- 
charged when  any  material  containing  uranium  is  brought 
near  it,  showing  that  the  air  all  about  this  material  is  an 
electrical  conductor. 

Fig.  183  represents  a  simple  form  of  aluminium-leaf  elec- 
troscope, the  leaves  assuming  the  position  indicated  at  B  when 
an  electric  charge  is  communicated  to  the  knob  A.  When  a 

substance  containing  ura- 
c  nium    (Fig.   184,    C)    is 

brought  near  the  knob, 
the  charge  is  rapidly  lost, 
and  the  leaves  collapse 
as  shown  at  B. 

The  discovery  of  ra- 
dium. Pitchblende  was 
found  to  be  4  times  as 
radioactive  as  uranium  itself,  which  suggested  that  possi- 
bly there  was  some  unknown  element  in  the  mineral  that 


FIG.  183.    An 
electroscope 


FIG.  184.    Discharging 
an  electroscope 


UKANIUM  AND  KADIUM 


381 


was  carried  over  into  the  uranium  salts  as  an  impurity  and 
was  responsible  for  the  radioactivity.  Accordingly  Monsieur 
and  Madame  Curie  worked  over  the  residues  from  a  very 
large  quantity  of  pitchblende,  and  obtained  a  minute  quan- 
tity of  the  chloride  of  a  new  element,  which  was  named 
radium.  This  chloride  is  about  3,000,000  times  as  active 
as  uranium. 

The  atomic  weight  of 
radium  is  226.4,  and  this 
weight,  as  well  as  all  of 
the  other  properties  of  the 
element  and  of  its  com- 
pounds, place  it  in  the  cal- 
cium family,  just  below 
barium.  The  metal  itself 
was  isolated  by  Madame 
Curie  (Fig.  185)  in  1910, 
and  is  very  similar  to 
barium. 

Quantity  of  radium  avail- 
able. Knowing  the  radio- 
activity of  both  uranium 
and  radium,  it  is  not  diffi- 
cult to  estimate  the  propor- 
tion of  radium  in  any  ore 

containing  uranium.  Estimates  of  this  kind  bring  to  light 
a  very  surprising  fact  —  the  proportion  of  radium  in  all 
classes  of  uranium  ores  is  very  constant,  and  is  about  1  part 
of  radium  in  2,940,000  parts  of  uranium. 

During  1914  there  was  produced  in  the  United  States 
87  tons  of  uranium  oxide,  and  at  the  ratio  just  given  this 
should  yield  29.4  g.  of  radium  chloride.  The  total  produc- 
tion from  foreign  ore  in  1912  was  less  than  4  g.  Probably 


FIG.  185.   Madame  Curie  (186 7~) 

Professor  of  Physics  in  the  University 
of  Paris 


382  FIRST  COUKSE  IN  CHEMISTRY 

the  total  quantity  produced  up  to  1914  is  less  than  50  g. 
The  present  price  of  the  chloride  is  about  $90,000  per  gram. 

Disintegration  of  radium.  The  extraordinary  fact  about 
radium  is  that  although  it  is  a  well-characterized  element, 
it  is  dowly  disintegrating.  In  this  process  it  is  resolved  into 
two  other  elements,  one  of  which  is  helium  and  the  other 
niton.  These  both  belong  in  periodic  Group  0  with  the 
inactive  gases  of  the  atmosphere.  Niton,  in  turn,  decom- 
poses into  helium  and  still  another  element  named  radium  A. 
Similar  decompositions  continue  through  a  number  of  stages, 
and  it  is  thought  that  the  final  product  is  lead. 

In  these  decompositions  two  distinct  kinds  of  particles  are 
shot  off  with  enormous  velocity :  (1)  the  one  kind,  called 
alpha  rays,  consists  of  helium  atoms  charged  positively  and 
moving  with  a  velocity  about  -^  that  of  light;  the  other, 
called  beta  rays,  consists  of  particles  not  more  than  ^y^-Q  of 
the  weight  of  a  hydrogen  atom  and  negatively  charged. 
These  are  usually  called  electrons,  and  their  initial  velocity 
is  nearly  that  of  light.  The  rate  at  which  this  decomposi- 
tion proceeds  cannot  be  changed  by  any  means  that  has 
yet  been  tried.  It  is  not  affected  by  very  high  temperature 
nor  by  the  nature  of  the  radium  compound. 

Origin  of  radium.  Radium  is  decomposing  at  a  rate  which 
places  its  average  life  at  2500  years,  yet  it  is  found  in  ores 
which  are  undoubtedly  much  older  than  this.  It  must  there- 
fore be  in  the  process  of  formation  from  some  other  ele- 
ment. Experiment  leaves  no  doubt  that  this  element  is 
uranium.  The  quantity  of  radium  so  constantly  present  in 
ores  of  uranium  simply  represents  the  equilibrium  between 
the  rate  at  which  uranium  disintegrates  and  that  at  which 
radium  disintegrates.  If  this  is  the  case,  it  is  clear  that 
we  can  never  hope  to  find  any  deposits  of  radium  richer 
than  those  afforded  by  uranium  ores. 


URANIUM  AND  EADIUM 


383 


Energy  of  radium.  During  the  decomposition  of  radium, 
niton,  and  the  succeeding  products,  a  very  great  deal  of 
energy  is  given  off.  Both  the  helium  atoms  and  the  elec- 
trons are  shot  off  with  very  high  kinetic  energy,  and  the 
radium  compound  is  kept  heated  by  the  heat  energy  set 
free.  It  is  estimated  that  1  gram  of  radium  hourly  evolves 
132  calories  of  heat.  From  this  value,  together  with  the 
average  life  period  (2500 
years),  it  is  easy  to  com- 
pute that  the  total  en- 
ergy given  off  by  a  gram 
of  radium  will  be  250,- 
000  times  the  heat  of 
combustion  of  a  gram 
of  carbon.  These  un- 
questioned facts  have 
thrown  a  great  deal  of 
doubt  upon  the  older 
estimates  of  the  age  of 
the  earth. 

Radium       and      the 

atomic  conception.     It  is       Fm-  186.    Total  amount  of  radium  bro- 

clear  that  the  atom  of      mide  t1-7,84*)  e?™(1  (rol"  30°  tons 

of  carnotite  (actual  size) 
radium,  as  well  as  that 

of  uranium,  must  have  a  very  elaborate  structure,  since 
helium,  electrons,  and  free  energy  are  formed  from  them. 
There  is  good  reason  for  thinking  that  all  atoms  have  a 
somewhat  similar  structure  —  a  little  like  a  miniature  solar 
system ;  but  other  atomic  systems  do  not  disintegrate  like 
the  ones  we  have  been  considering. 

Radium  and  medicine.  The  rays  emitted  from  radium, 
niton,  and  other  radioactive  elements  produce  many  chemi- 
cal and  physiological  effects.  They  disintegrate  glass,  water, 


384  FIKST  COUESE  IN  CHEMISTEY 

and  many  other  substances.  They  produce  severe  burns 
upon  the  skin,  like  those  of  X  rays.  They  kill  bacteria 
and  other  microorganisms. 

This  latter  property  has  led  to  the  hope  that  exposure  to 
the  radiations  of  radium  compounds  might  prove  to  be  of 
assistance  in  effecting  a  cure  for  some  diseases  of  the  skin 
and  for  cancer.  It  is  not  possible  as  yet  to  say  to  what  ex- 
tent these  hopes  will  be  realized.  Certain  forms  of  cancer 
have  apparently  been  cured  in  this  way. 

Radioactive  thorium.  The  rare  element  thorium  exhibits 
properties  very  similar  to  those  of  uranium.  It  gives  rise 
to  the  same  kind  of  a  series  of  radioactive  elements  by 
successive  decomposition,  producing  the  same  varieties  of 
radiation  as  the  other  series.  Uranium  and  thorium  are 
the  elements  of  greatest  atomic  weight,  and  no  others  are 
known  to  possess  similar  properties.  This  suggests  the 
idea  that  possibly  elements  of  still  higher  atomic  weight 
may  have  existed  at  some  time,  but  that  they  have  disin- 
tegrated to  form  elements  of  smaller  atomic  weight  which 
are  not  radioactive. 

EXERCISES 

1.  When  was  uranium  discovered,  and  how  did  it  get  its  name 
(see  encyclopedia)  ? 

2.  For  whom  was  carnotite  named  (see  dictionary)? 

3.  How  is  an  electroscope  charged  (see  physics)? 

4.  What  is  the  meaning  of  alpha  and  beta? 

5.  What  is  the  velocity  of  light  (see  physics)? 

6.  I  low  did  thorium  get  its  name  ?   For  what  is  the  element  used  ? 

TOPICS  FOR  THEMES 

The  discovery  of  radium  (McPherson  and  Henderson,  Course  in 
General  Chemistry). 

Radium  ores  of  the  United  States  (Bureau  of  Mines,  Washing- 
ton, D.C.). 

Radium  in  medicine.    (See  current  magazines.) 


CHAPTER  XL 
MANGANESE  AND  CHROMIUM 


NAME 

SVMHOL 

ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POIXT 

FORMULAS  OF  ACIDS 

Manganese 
Chromium 

Mn 

Cr 

54.93 
52. 

7.39 
6.50 

1260° 
1520° 

II2MnO4  and  IIMnO4 
II2004  and  H2O2O7 

MANGANESE 

Occurrence.  Manganese  is  found  in  nature  chiefly  as  the 
dioxide  MnO2,  called  pyrolmite.  In  smaller  amounts  it 
occurs  as  the  oxides  Mn  C)  and  Mn  O  ,  and  as  the  carbon- 

2       o  o       4 

ate  MnCOg.    Some  iron  ores  also  contain  manganese. 

Preparation  and  properties.  The  metal  is  difficult  to  pre- 
pare in  pure  condition,  and  has  no  commercial  applications. 
It  somewhat  resembles  iron  in  appearance,  but  is  harder, 
more  fusible,  and  more  readily  acted  upon  by  air  and 
moisture.  Acids  readily  dissolve  it,  forming  manganous 
salts.  An  alloy  of  manganese  and  iron,  called  ferroman- 
ganese,  is  made  by  reducing  a  mixture  of  oxides  of  the 
two  metals,  and  is  used  in  the  steel  industry. 

Manganese  dioxide  (pyrolusite)  (Mn02).  This  substance 
is  the  ore  from  which  all  other  compounds  of  manganese 
are  made.  It  is  a  hard,  brittle,  black  substance  which  is 
valuable  as  an  oxidizing  agent.  It  will  be  recalled  that  it 
has  been  used  in  the  preparation  of  chlorine  and  oxygen, 
and  in  decolorizing  glass  which  contains  iron.  At  present 
its  chief  use  is  in  the  manufacture  of  ferromanganese. 

385 


886  FIKST  COURSE  IK  CHEMISTKY 

Manganous  salts.  Manganese  acts  as  a  bivalent  metal, 
forming  a  series  of  salts  with  all  the  common  acids.  The 
chloride  and  the  sulfate  may  be  prepared  by  heating  the 
dioxide  with  hydrochloric  acid  and  sulfuric  acid  respec- 
tively. The  sulfide,  carbonate,  and  hydroxide,  being  insol- 
uble, may  be  prepared  from  a  solution  of  the  chloride 
or  sulfate  by  precipitation  with  the  appropriate  reagents. 
Most  of  the  manganous  salts  are  pink.  They  not  only  have 
formulas  similar  to  the  ferrous  salts  but  resemble  them  in 
many  of  their  chemical  properties. 

The  formulas  of  some  of  these  salts  are  as  follows : 

Manganous  chloride        MiiCl0  •  4  IT2O 

Manganous  sulfide MuS 

Manganous  sulfate MnSO4  •  4  H2O 

Manganous  carbonate MnCO3 

Potassium  permanganate  (KMnOJ.  When  manganese 
dioxide  is  fused  with  potassium  hydroxide  and  an  oxidizing 
agent,  and  the  resulting  mass  is  extracted  with  water,  a 
deep-green  solution  is  obtained.  This  color  changes  to  a 
very  intense  reddish  purple  when  carbon  dioxide  is  passed 
into  the  solution,  and  upon  evaporating,  purple-black  crys- 
tals of  potassium  permanganate  (KMnO4)  are  formed.  This 
is  the  potassium  salt  of  permanganic  acid  (HMnO4).  The 
free  acid  gives  a  solution  of  the  same  intense  color  as  do 
all  permanganates. 

Oxidizing  properties  of  potassium  permanganate.  Potas- 
sium permanganate  is  remarkable  for  its  strong  oxidizing 
properties,  especially  in  the  presence  of  an  acid.  When  sul- 
furic acid  is  present,  the  reaction  takes  place  in  such  a  way 
that  both  the  potassium  and  the  manganese  are  changed  into  sul- 
fates,  with  the  liberation  of  oxygen,  as  shown  in  the  equation 

•1  K  MnO4  +  3  H2SO4 >-  K2SO4  +  2  MnSO4  +  3  H2O  +  5  [O] 


MANGANESE  AND  CHROMIUM  387 

Under  ordinary  conditions,  however,  the  reaction  does  not 
take  place  except  in  the  presence  of  a  third  substance 
which  is  capable  of  oxidation.  The  oxygen  is  not  given 
off  in  the  free  state,  but  is  used  up  in  effecting  oxidation. 
This  is  indicated  by  inclosing  its  symbol  in  brackets. 

CHROMIUM 

Occurrence.  The  ore  from  which  all  chromium  com- 
pounds are  made  is  chromite,  or  chrome  iron  ore  (FeCr2O4). 
This  is  found  most  abundantly  in  New  Caledonia,  Greece, 
and  California.  The  element  also  occurs  in  small  quantities 
in  many  other  minerals. 

Preparation  and  properties.  Chromium,  like  manganese, 
is  very  hard  to  reduce  from  its  ores,  owing  to  its  great 
affinity  for  oxygen.  It  can  be  produced  by  the  Gold- 
schmidt  method  (p.  325).  The  pure  metal  has  no  com- 
mercial uses,  but  an  alloy  of  chromium  with  iron,  called 
f err o chromium,  is  used  in  the  steel  industry.  Chromium 
is  a  very  hard  metal  of  about  the  same  density  as  iron. 

Chromic  hydroxide  (Cr(OH)3).  This  substance,  being 
insoluble,  can  be  obtained  by  precipitating  a  solution  of 
the  chloride  or  the  sulfate  with  a  soluble  hydroxide.  It  is 
a  greenish  substance  which,  like  aluminium  hydroxide,  dis- 
solves in  excess  of  alkalies.  Both  the  hydroxide  and  the 
corresponding  oxide,  Cr0Og,  are  used  as  green  pigments. 

Chromic  salts.  Chromium  acts  as  a  trivalent  metal,  form- 
ing a  series  of  salts  resembling  ferric  salts  in  formula.  Most 
of  them  are  green  or  violet  in  color.  The  most  important 
of  the  chromic  series  are  the  following: 

Chromic  chloride CrCl8  •  6  H2O 

Chromic  sulfate Cr2(SO4)3 

Potassium  chrome  alum KCr(SO4)2  •  12  H2O 


388  FIRST  COURSE  IX  CHEMISTRY 

A  number  of  the  salts  of  chromium  are  used  in  the  dye- 
ing industry,  for  they  hydrolyze  like  aluminium  salts,  and 
the  hydroxide  forms  a  good  mordant. 

Potassium  chromate  (K2Cr04).  When  a  chromium  com- 
pound is  fused  with  potassium  hydroxide  and  an  oxidizing 
agent,  potassium  chromate  (K2CrO4)  is  formed.  With  chro- 
mium hydroxide  the  equation  is  as  follows : 

2  Cr(OH)8  +  4  KOH  +  3  [O] >•  2  K2CrO4  +  5  H2O 

Properties  of  chromates.  The  chromates  are  salts  of  the 
unstable  chromic  acid  (H2CrO4),  and  as  a  rule  are  yellow  in 
color.  Most  of  the  chromates  are  insoluble,  and  can  be  pre- 
pared from  the  soluble  potassium  salt  by  precipitation.  In 
the  case  of  lead  chromate  the  equation  is  as  follows : 

Pb(N03)2  +  K2Cr04 ^  PbCr04  +  2  KNO8 

Lead  chromate  (chrome  yellow)  and  barium  chromate  are 
used  as  yellow  pigments. 

Potassium  dichromate  (K2Cr207).  When  potassium  chro- 
mate is  treated  with  sulfuric  acid,  the  potassium  salt  of 
dichromic  acid  (H2Cr2O7)  is  formed : 

2  K,Cr04  +  H2S04  — f  K.Cr.0,  +  K2SO4  +  H2O 

This  is  the  best-known  dichromate,  and  is  the  most  familiar 
chromium  compound.  It  forms  large  crystals  of  a  brilliant- 
red  color,  and  is  rather  sparingly  soluble  in  water.  Potassium 
dichromate  finds  use  in  many  industries  as  an  oxidizing 
agent,  especially  in  the  preparation  of  organic  substances, 
such  as  the  dye  alizarin,  and  in  the  construction  of  several 
kinds  of  electrical  batteries. 

Leather.  Leather  is  made  from  the  skins  of  various  animals, 
the  processes  employed  being  such  as  to  make  the  material 
pliant,  impervious  to  water,  and  not  subject  to  putrefaction. 


MANGANESE  AND  CHROMIUM  389 

The  hair  is  removed  by  treatment  with  milk  of  lime,  and  the 
skin  is  then  washed  with  dilute  acid  to  neutralize  the  lime 
and  cause  the  skin  to  swell.  It  is  then  soaked  in  some  tan- 
ning solution,  the  active  reagent  of  which  combines  with  the 
nitrogenous  constituents  of  the  skin  and  changes  them  into 
substances  which  do  not  decay.  Formerly  the  bark  of  hemlock 
and  oak  trees,  which  contain  tannin,  was  exclusively  used  for 
this  purpose,  but  more  recently  potassium  dichromate  is  often 
used,  as  by  this  method  the  time  required  for  tanning  is  much 
shortened.  Finally,  the  tanned  skin  is  treated  with  oil  to  make 
it  soft  and  waterproof. 

Oxidizing  action  of  potassium  dichromate.  When  a  dilute 
solution  of  potassium  dichromate  is  treated  with  sulfuric 
acid,  no  reaction  apparently  takes  place.  However,  if  there 
is  present  a  third  substance  capable  of  oxidation,  the  dichro- 
mate gives  up  a  portion  of  its  oxygen  to  this  substance,  and 
both  the  potassium  and  the  chromium  are  converted  into 
sulfates.  The  oxidation  of  ferrous  sulfate  by  potassium 
dichromate  is  a  good  illustration,  the  reaction  being  repre- 
sented in  two  steps  : 


6  FeS04  +  3  H2S04  +  3  [O]  --  +  3  Fea(SO4)8  +  3  H2O 

This  reaction  is  often  employed  in  the  analysis  of  iron  in 
iron  ores. 


EXERCISES 


1.    How  does  pyrolusite  effect  the  decolorizing  of  glass  contain- 


ng  ron  ? 


2.  Write  the  equations  for  the  preparation  of  manganous  nitrate, 
manganous  carbonate,  and  manganous  hydroxide. 

3.  Write  the  equations  representing  the  reactions  which  take 
place  when  ferrous  sulfate  is  oxidized  to  ferric  sulfate  by  potassium 
permanganate  in  the  presence  of  sulfuric  acid. 


390  FIRST  COURSE  IN  CHEMISTRY 

4.  Potassium    permanganate   is   sometimes    injected    as    a    cure 
around  the  wound  caused  by  a  rattlesnake  bite.    How  would  you 
suppose  it  acts? 

5.  Where  is  New  Caledonia  ? 

6.  Why  do  all  permanganates  have  the  same  color  in  solution? 

7.  10  g.  of  iron  was  dissolved  in   sulfuric  acid  and  oxidized  to 
ferric  sulfate  by  potassium  permanganate.    What  weight  of  the  per- 
manganate was  required?   Ans.  5.66  g. 

8.  Potassium  chromate  oxidizes  hydrochloric  acid,  forming  chlo- 
rine.   Write  the  complete  equation. 

9.  20  Ib.  of  ferrochromium  containing  40%  chromium  was  added 
to   a  ton   of   steel.   What  per  cent  of  chromium  did  the  product 
contain?    Ans.  O.o98. 

TOPICS  FOR  THEMES 

The  alums  (see  encyclopedia). 

How  sole  leather  is  made  (Lassar-Cohn,  Chemistry  in  Daily  Life). 


CHAPTER  XLI 
PLATINUM  AND  GOLD 


NAME 

SYMBOL 

ATOMIC 
WEIGHT 

DENSITY' 

HIGHEST 
OXIDE 

HIGHEST 
CHLORIDE 

MELTING 
POINT 

Platinum 
Gold    .     . 

Pt 

Au 

195.2 
197.2 

21.50 
19.32 

PtO2 

Au203 

PtCl4 
AuCl8 

1755° 
1062° 

PLATINUM 

Occurrence.  About  90  per  cent  of  the  platinum  of  com- 
merce comes  from  Russia,  small  amounts  being  produced 
in  California,  Brazil,  and  Colombia.  Like  gold,  it  usually 
occurs  in  metallic  grains  in  heavy  sands,  but  it  has  recently 
been  found  in  eastern  Germany  in  combined  form. 

Preparation.  Native  platinum  is  usually  alloyed  with 
gold  and  other  rare  metals.  To  separate  the  platinum,  the 
alloy  is  dissolved  in  aqua  regia,  which  converts  the  platinum 
into  cliloroplatinic  acid  (HoPtCl6)^  Ammonium  chloride  is 
then  added,  which  precipitates  the  platinum  as  insoluble 
ammonium  chloroplatinate : 

H2PtCl6+  2  NH4C1  — >•  (NH4)2PtCl6+  2  HC1 

On  heating  the  ammonium  compound,  it  is  decomposed, 
leaving  the  platinum  as  a  powdery  metallic  mass  known  as 
platinum  sponge.  This  may  be  melted  into  an  ingot  in  an  elec- 
tric furnace  and  rolled  or  hammered  into  any  desired  shape. 

391 


392  FIRST  COUBSE  IN  CHEMISTRY 

Properties.  Platinum  is  a  grayish-white  metal  of  high 
luster,  and  is  very  malleable  and  ductile.  It  melts  in  the 
oxyhydrogen  blowpipe  and  in  the  electric  furnace,  and 
is  a  good  conductor  of  electricity.  In  finely  divided  form 
it  has  the  ability  to  absorb,  or  occlude,  gases,  especially 
oxygen  and  hydrogen.  These  gases,  when  occluded,  are  in 
a  very  active  condition  resembling  the  nascent  state,  and 
can  combine  with  each  other  at  ordinary  temperatures.  A 
jet  of  hydrogen  or  coal  gas  directed  upon  spongy  platinum 
quickly  ignites. 

Platinum  as  a  catalytic  agent.  Platinum  is  remarkable  for 
its  property  of  acting  as  a  catalytic  agent  in  a  large  number  of 
chemical  reactions,  and  mention  has  been  made  of  this  use  of 
the  metal  in  connection  with  the  manufacture  of  sulfuric  acid. 
When  desired  for  this  purpose  some  porous  or  fibrous  substance, 
such  as  asbestos,  is  soaked  in  a  solution  of  chloroplatinic  acid 
and  then  ignited.  The  platinum  compound  is  decomposed  and 
the  platinum  deposited  in  very  finely  divided  form.  Asbestos 
prepared  in  this  way  is  called  platinized  asbestos.  The  catalytic 
action  seems  to  be  in  part  connected  with  the  property  of  ab- 
sorbing gases  and  rendering  them  nascent.  Some  other  metals 
possess  this  same  power,  notably  palladium,  which  is  remarkable 
for  its  ability  to  absorb  hydrogen. 

Chemical  conduct.  Platinum  is  a  very  inactive  element 
chemically,  and  is  not  attacked  by  any  of  the  common 
acids.  Aqua  regia  slowly  dissolves  it,  and  it  is  also  at- 
tacked by  fused  alkalies.  It  combines  at  higher  tempera- 
tures with  carbon  and  phosphorus,  and  forms  alloys  with 
many  metals.  It  is  readily  attacked  by  chlorine  but  not  by 
oxidizing  agents. 

Applications.  Platinum  is  very  valuable  as  a  material  for 
the  manufacture  of  chemical  utensils  which  are  required  to 
stand  a  high  temperature  or  the  action  of  strong  reagents. 


PLATINUM  AND  GOLD  393 

Platinum  crucibles,  dishes,  forceps,  electrodes,  and  similar 
articles  (Fig.  187)  are  indispensable  in  the  chemical  labora- 
tory. In  the  industries  platinum  is  used  for  such  purposes 
as  the  manufacture  of  pans  for  evaporating  sulfuric  acid, 
and  wires  for  sealing  through  incandescent-light  bulbs ;  and 
also  as  a  catalytic  material  in  a  number  of  reactions,  and 
for  making  a  great  variety  of  instruments.  A  large  fraction 
of  the  annual  production  is  used  for  jewelry.  Unfortunately 
the  supply  of  the  metal  is  very  limited,  and  its  cost  is  stead- 
ily advancing,  so  that  it  is  now  more  valuable  than  gold. 


FIG.  187.    Some  laboratory  utensils  made  of  platinum 

Chloroplatinic  acid  (H2PtCl6).  When  platinum  is  dis- 
solved in  aqua  regia  and  the  solution  is  evaporated  to 
dryness,  orange-colored  crystals  of  chloroplatinic  acid 
(H2PtCl6)  are  obtained.  The  potassium  and  ammonium 
salts  of  this  acid  are  nearly  insoluble  in  water  and  alcohol. 

GOLD 

Occurrence.  Gold  has  been  found  in  many  localities,  the 
best  known  being  South  Africa,  Australia,  Russia,  and  the 
United  States.  In  this  country  it  is  found  in  Alaska  and  in 
nearly  half  of  the  states  of  the  union,  notably  in  California, 
Colorado,  and  Nevada.  It  is  usually  found  in  the  native 


394  FIRST  COUKSE  IN  CHEMISTRY 

condition,  frequently  alloyed  with  silver ;  in  combination 
it  is  sometimes  found  as  telluride  (AuTe2).  The  United 
States  produces  over  one  fifth  of  the  world's  annual  output. 

Extraction.  The  extraction  of  gold  is  accomplished  in  a 
number  of  ways,  according  to  the  character  of  the  deposit. 
In  placer  mining  the  gold-bearing  sand  is  washed  by  a  cur- 
rent of  water  which  is  so  regulated  that  particles  of  light 
weight  are  swept  away,  while  the  heavier  gold  is  obtained 
as  a  sediment.  In  hydraulic  mining  the  earth  and  sand  are 
swept  into  sluices  by  powerful  streams  of  water  operated  by 
pumps.  In  quartz  mining  the  quartz  is  stamped  to  powder 
and  is  then  washed  over  copper  plates,  the  surfaces  of  which 
have  been  amalgamated.  The  particles  of  gold  stick  to  the 
mercury  or  dissolve  in  it,  the  gold  being  recovered  by  dis- 
tillation. In  other  cases,  especially  when  the  gold  is  in  very 
fine  powder  or  in  chemical  combination,  chemical  reactions 
are  employed.  In  the  cyanide  process  the  gold-bearing  mate- 
rial is  treated  with  a  dilute  solution  of  potassium  cyanide, 
with  free  access  of  air.  The  gold  dissolves  to  form  a  com- 
plex cyanide,  from  which  it  can  be  precipitated  by  metallic 
zinc  or  by  electrolysis. 

In  the  chlorination  process  the  ore  is  treated  with  chlorine, 
which  converts  the  gold  into  the  soluble  trichloride  AuCl3. 
It  is  recovered  from  this  solution  by  suitable  precipitants. 
The  treatment  of  lead  and  silver  ores  containing  gold,  as 
well  as  the  separation  of  gold  from  silver,  has  already  been 
described  (p.  373). 

Properties.  Gold  is  a  very  heavy  bright-yellow  metal, 
exceedingly  malleable  and  ductile,  and  a  good  conductor 
of  electricity.  Its  melting  point  (1062°)  is  much  below 
that  of  platinum.  It  is  quite  soft,  and  is  usually  alloyed 
with  copper  or  silver  to  give  it  the  hardness  required  for 
most  practical  uses.  The  degree  of  fineness  is  expressed  in 


PLATINUM  AND  GOLD  395 

terms  of  carats,  pure  gold  being  24  carats;  the  gold  used 
for  jewelry  is  usually  18  carats,  18  parts  being  gold  and 
6  parts  copper  or  silver.  Gold  coinage  is  90  per  cent  gold 
and  10  per  cent  copper. 

Chemical  conduct.  Gold  is  not  attacked  by  any  one  of 
the  common  acids ;  aqua  regia  easily  dissolves  it,  forming 
cJilorauric  acid  (HAuCl4).  Fused  alkalies  also  attack  it. 
Most  oxidizing  agents  are  without  action  upon  it,  and  in 
general  it  is  not  an  active  element. 

EXERCISES 

1 .  What  is  the  derivation  of  the  word  platinum  ? 

2.  The  "platinum  chloride"  of  the  laboratory  is  made  by  dis- 
solving platinum  in  aqua  regia.    What  is  this  compound? 

3.  What  properties  would  wire  need  to  have  in  order  to  be 
adapted  to  sealing  through  electric-light  bulbs  ? 

4.  Since  gold  is  cheaper,  why  is  it  not  used  instead  of  platinum 
for  laboratory  utensils  ? 

5.  How  could  gold  be  precipitated  from  the  chloride? 

TOPICS  FOR  THEMES 

Discovery  of  gold  in  California  (see  encyclopedia). 

The  Alaska  gold  fields  (U.  S.  Geological  Survey, Washington,  D.C.). 

The  uses  of  platinum  (see  encyclopedia). 


CHAPTER  XLII 
SOME  APPLICATIONS  OF  RARER  ELEMENTS 

Rarer  elements.  A  large  number  of  elements  are  known 
which  have  not  been  described  in  the  foregoing  pages  be- 
cause an  acquaintance  with  them  is  not  at  all  necessary  for 
an  understanding  of  the  principles  of  chemistry. 

Some  of  these,  while  comparatively  rare,  could  be  pro- 
duced in  considerable  quantities  if  there  were  any  commer- 
cial use  for  them.  A  good  example  is  tellurium,  an  element 
in  the  sulfur  family  obtained  as  a  by-product  in  copper 
refining.  Others  of  these  elements  are  so  rare  that  the  cost 
of  production  is  prohibitive,  even  though  they  have  very 
useful  properties. 

Application  in  the  industries.  Some  of  these  less  famil- 
iar elements  or  their  compounds  have  properties  which 
make  them  valuable  for  special  purposes,  and  mention  of 
a  few  of  these  applications  will  be  of  interest. 

The  rare  earths  constitute  a  group  of  about  sixteen 
elements,  all  trivalent  and  resembling  aluminium  in  a  gen- 
eral way.  They  are  very  difficult  to  separate  from  each 
other  and  always  occur  together  in  nature.  Very  large 
quantities  of  a  mixture  of  them  accumulate  in  the  extrac- 
tion of  thorium  from  monazite  sand  (p.  213).  The  only  one 
whose  compounds  are  obtained  pure  rather  easily  is  cerium. 
Compounds  of  cerium  are  used  as  mordants,  as  catalytic 
agents,  and  in  medicine  and  photography..  An  alloy  of 
cerium  with  iron  is  used  as  a  gas  or  cigar  lighter,  since  it 
gives  off  a  stream  of  sparks  when  scratched  by  hard  iron. 

396 


APPLICATIONS  OF  KAKEK  ELEMENTS       397 


Thorium  oxide,  mixed  with  1  per  cent  of  cerium  oxide, 
constitutes  the  material  of  which  most  gas  mantles  are 
made  (Fig.  101). 

Titanium  in  the  silicon  family  is  not  a  very  rare  element, 
occurring  chiefly  as  the  oxide  TiO0,  called  rutile,  and  as  a 
constituent  of  certain  iron  ores  (ilmenite).  Large  quantities 
of  nearly  pure  titanium  or  of  ferrotitanium  are  used  in 
making  steel  rails  designed  to  stand  very  heavy  wear 
(railway  curves  and  terminals).  Titanium  oxide  is  also 
incorporated  in  electric-arc  carbons  {flam- 
ing arc).  Carbons  thus  made  give  a  more 
diffused  and  efficient  light  than  those  made 
from  pure  carbon.  The  oxide  is  also  used 
to  impart  a  yellow  color  to  porcelain  and 
to  artificial  teeth. 

Vanadium  also  occurs  in  considerable 
quantities  in  carnotite  (p.  379)  and  in  cer- 
tain sulfides  found  in  Peru.  It  is  found 
as  oxide  in  the  ash  of  nearly  all  anthracite 
coal.  Ferrovanadium,  like  ferrotitanium, 
is  used  in  producing  special  grades  of  steel, 
particularly  when  great  toughness  is  desired  (automobile 
parts).  Its  compounds  are  used  as  photographic  developers, 
as  catalytic  reagents  in  the  dye  industry  (aniline  black),  as 
coloring  materials  in  glass,  and  as  mordants. 

Molybdenum  compounds  are  used  in  coloring  pottery  and 
in  dyeing  silk,  wool,  and  leather. 

Tungsten  compounds  are  produced  in  fairly  large  quan- 
tities. It  has  been  found  possible  to  draw  the  metal  into 
very  fine  wire  (0.3  mm.),  which  is  now  extensively  used 
instead  of  carbon  as  a  filament  for  incandescent  lamps 
(Fig.  188).  Its  melting  point  is  very  high  (3000°),  and 
the  consumption  of  electrical  energy  for  a  given  candle 


FIG.  188.  A  tung- 
sten lamp 


FIRST  COURSE  IN  CHEMISTRY 

power  is  so  low  that  the  lamp  is  about  three  times  as 
efficient  as  the  older  (carbon)  lamp.  The  metal  is  rapidly 
replacing  platinum  for  electrical  contacts  in  switches,  tele- 
phone jacks,  and  automobile  vibrators.  Ferrotungsten  is 
used  in  making  steel  designed  for  lathe  tools,  since  such 
steel  can  be  heated  to  a  red  glow  without  losing  temper. 

Compounds  of  tungsten  are  used  for  making  fireproof 
cloth,  and  pigments  for  paints  and  pottery,  and  as  mordants. 

Selenium,  an  element  in  the  sulfur  family,  is  obtained 
as  a  by-product  in  refining  copper.  It  is  a  nonconductor 
of  electricity  when  in  the  dark,  but  becomes  a  fairly  good 
conductor  when  exposed  to  "light.  This  has  led  to  its  use  in 
automatic  fire  alarms  and  for  regulating  automatic  gas 
buoys  at  sea.  Added  to  glass  it  produces  a  fine  red  color, 
such  glass  being  used  for  railway  lanterns.  It  is  also  used 
to  produce  red  enamels. 

Iridium  gives  a  very  hard  alloy  with  platinum,  used  for  pen 
points,  compass  bearings,  and  standard  weights  and  measures. 

Palladium  is  only  about  half  as  heavy  as  platinum,  melts 
much  lower,  and  is  harder.  It  is  used  as  a  solder  for  platinum, 
for  making  graduated  scales  in  scientific  instruments,  and  as  a 
substitute  for  platinum  in  jewelry.  In  the  form  of  a  powder  it 
is  a  remarkably  active  catalytic  agent. 


APPENDIX 


THEME-WRITING  IN  CHEMISTRY 

Following  each  chapter  will  be  found  one  or  more  topics  for 
themes.  Much  interest  will  be  added  to  the  study  of  chemistry 
if  the  pupil  is  encouraged  to  look  up  for  himself  some  of  the 
details  of  the  lives  of  the  most  illustrious  chemists  and  to  find 
out  by  reading,  by  consultation,  and  by  observation  something 
more  about  the  applications  of  chemistry  in  the  arts  and  indus- 
tries than  can  be  described  in  the  pages  of  a  brief  text.  These 
topics  are  intended  to  be  merely  suggestive.  The  teacher 
should  add  others,  especially  those  concerning  which  the 
student  will  be  able  to  obtain  first-hand  information. 

It  is  suggested  that  there  should  be  close  cooperation  between 
the  teacher  of  chemistry  and  the  teacher  of  English.  The  same 
theme  may  be  read  in  the  chemistry  class  for  its  contents  and 
in  the  English  class  for  its  form.  This  ought  to  add  interest 
to  both  classes  without  increasing  the  work  of  the  student  and 
at  the  same  time  will  emphasize  the  important  fact  that  the 
writing  of  good  English  should  be  not  merely  an  exercise  for 
the  English  classroom  but  a  habit  in  all  written  discourse. 

If  practicable,  the  student  should  visit  actual  plants  in  opera- 
tion, and  theme  topics  should  be  chosen  to  meet  local  conditions. 
Many  large  firms  can  supply  printed  descriptions  of  their  proc- 
esses and  products,  and  a  courteous  letter  of  inquiry  is  almost 
certain  to  bring  a  courteous  reply.  It  would  be  well  for  the 
teacher  to  supervise  and  criticize  such  a  letter  to  see  that  it  is 
in  proper  form  and  spirit.  Printed  matter  secured  in  this  way 
should  be  preserved  in  available  form  as  the  property  of  the 
school.  Desired  information  may  often  be  had  from  a  local 

399 


400  FIRST  COUESE  IN  CHEMISTRY 

dealer  or  agent,  and  in  almost  every  community  there  are 
well-informed  persons  who  would  gladly  give  some  time  to  an 
inquiring  student.  Training  in  the  gathering  of  trustworthy 
information  is  far  more  important  than  the  information  itself. 
The  course  in  chemistry  affords  unusual  opportunities  for  such 
training,  and  it  is  worth  some  effort  to  secure  it. 

Much  that  is  interesting  in  regard  to  the  applications  of 
chemistry  will  be  found  in  the  files  of  various  periodicals,  such 
as  the  Scientific  American,  the  World's  Work,  and  School  Science 
and  Mathematics.  The  bulletins  published  by  the  United 
States  Department  of  Agriculture  often  contain  information 
of  great  value  to  the  student  of  chemistry  and  they  cost  but  a 
few  cents.  A  list  of  available  bulletins  may  be  obtained  by 
addressing  the  Department  of  Agriculture,  Washington,  D.  C. 

A  list  of  books  is  appended  which  will  be  found  helpful.  The 
list  might  easily  be  greatly  extended,  but  is  purposely  brief  so 
as  to  bring  it  within  the  reach  of  almost  any  high  school.  In 
connection  with  the  topics  for  themes  frequent  reference  is 
made  to  the  encyclopedia,  partly  because  it  is  almost  sure  to  be 
at  hand  and  partly  because  it  is  desirable  to  teach  boys  and 
girls  to  use  it  freely.  Almost  any  public  library  will  be  glad 
to  furnish  a  list  of  its  books  relating  to  chemical  topics,  and 
such  a  list  should  be  posted  in  the  chemical  laboratory. 

LIST  OF  SUPPLEMENTARY  BOOKS 

ALLYN.    Elementary  Applied  Chemistry.    Ginn  and  Company. 
BAILEY.    A  Textbook  of   Sanitary  and   Applied   Chemistry.    The 

Macmillan  Company. 

BIRD.   Modern  Science  Reader.    The  Macmillan  Company. 
BLOXAM.  Inorganic  and  Organic  Chemistry.  P.  Blakiston's  Son  &  Co. 
DAVY.    The  Elementary  Nature  of  Chlorine.     The  University  of 

Chicago  Press. 

DUNCAN.    The  Chemistry  of  Commerce.    Harper  &  Brothers. 
DUNCAN.    The  New  Knowledge.    Harper  &  Brothers. 
FARADAY.    The  Liquefaction  of  Gases.    The  University  of  Chicago 

Press. 


APPENDIX  401 

FREAR.    Breakfast  Foods,  Bulletin  162,  Dairy  and  Food  Division, 

Department  of  Agriculture,  Harrisburg,  Pa. 
LASSAR-COHN    (translated    by    Muir).     Chemistry   in    Daily    Life. 

II.  Grevel  Co.,  London. 
McPnEKSON  and  HENDERSON.    A   Course   in    General    Chemistry. 

Ginn  and  Company. 
MARTIN.    Triumphs  and  Wonders  of  Modern  Chemistry.    D.  Van 

Nostrand  Company. 

MUIR.    The  Story  of  Alchemy.    D.  Appleton  and  Company. 
PRIESTLEY.    The  Discovery  of  Oxygen,  Part  I.    The  University  of 

Chicago  Press. 
ROGERS  and  AUBERT.    Industrial    Chemistry.    D.  Van   Nostrand 

Company. 
ROSCOE  and  SCHORLEMMER.    Inorganic  Chemistry,  Vols.  I  and  II. 

The  Macmillan  Company, 
SCHEELE.    The  Discovery  of  Oxygen,  Part  II.    The  University  of 

Chicago  Press. 
SCHEELE.   The  Early  History  of  Chlorine.  The  University  of  Chicago 

Press. 
SHERMAN.     Chemistry   of    Food   and   Nutrition.     The    Macmillan 

Company. 

STEWART.    Chemistry  and  its  Borderland.    Longmans,  Green,  &  Co. 
THORPE.    Essays  in  Historical  Chemistry.  The  Macmillan  Company. 
VENABLE.    A  Short  History  of  Chemistry.    D.  C.  Heath  &  Co. 
WILEY.    Foods  and  their  Adulteration.    P.  Blakiston's  Son  &  Co. 
United  States  Department  of  Agriculture  :  (1)  Composition  of  Foods, 
Bulletin  28,  Office  of  Experiment  Stations ;    (2)  Nutritive  Value 
of  Foods,  Farmers'  Bulletin  142]   (3)  Some  Forms  of  Food  Adul- 
teration  and   Simple    Methods  for  their  Detection,  Bulletin   100, 
Bureau  of  Chemistry;  (4)  Industrial  Alcohol,  Farmers' Bulletins  268 
and  269  ;   (5)  Household  Tests  for  the  Detection  of  Oleomargarine 
and  Renovated  Butter,  Farmers'  Bulletin  231  ;   (6)  The  Use  of  Milk 
as  Food,  Farmers'  Bulletin  363 ;   (7)  Canned  Fruits,  Preserves,  and 
Jellies,  Farmers'  Bulletin  203.     (Send  to  the  Department  of  Agri- 
culture, Washington,   D.C.,  for  list  of  available  bulletins,  and 
select  such  as  may  be  of  interest.)* 


402 


FIRST  COURSE  IN  CHEMISTRY 


TENSION  OF  AQUEOUS  VAPOR  AT  VARIOUS  TEMPERATURES, 
EXPRESSED  IN  MILLIMETERS  OF  MERCURY 


TEMPERATUKE 
0°  .... 
16°  .  .  .  . 
17°  .... 
18°  .  .  .  . 
19°  .  .  .  . 
20°  . 


PRESSURE  TEMPERATURE  PRESSURE 

.     .  4.6  21° 18.62 

.     .  13.62  22° 19.79 

.     .  14.4  23° 21.02 

.     .  15.46  24° 22.32 

.     .  16.45  25° 23.69 

.  17.51  100°  760.00 


WEIGHT  IN  GRAMS  OF  1  LITER  OF  VARIOUS  GASES  UNDER  STAND- 
ARD  CONDITIONS,   AND   BOILING   POINTS   UNDER   PRESSURE    OF 
760  MILLIMETERS 


WEIGHT 
NAME          OF  1  LITER 

Acetylene  .     .     .1.1621 
Air    1.2928 

BOILING 

POINT 

-83.8° 

-33.5° 
-186.0° 
-78.2° 
—  190.0° 
-33.6° 
-268.7° 
-252.7° 

WEIGHT 
NAME          OF  1  LITER 

Hydrogen  chloride  1.6398 
Hydrogen  fluoride  0.893 
Hydrogen  sulfide    1.5392 
Methane     .     .     .0.7168 
Nitric  oxide     .     .  1.3402 
Nitrogen     .     .     .1.2507 
Nitrous  oxide       .  1.9777 
Oxygen       ...  1.4290 
Sulfur  dioxide     .  2.9266 

BOILING 

POINT 
-82.9° 
+  19.4° 
-61.6° 
-164.0° 
-153.0° 
-195.7° 
—  89.8° 
-182.9° 
—  10.1° 

Ammonia  .     .     . 
Argon    .... 
Carbon  dioxide    . 
Carbon  monoxide 
Chlorine     .     .     . 
Helium  .... 
Hydrogen  . 

0 
1 

1 
1 
3 
0 
0 

.7708 
.7809 
.9768 
.2504 
.1674 
.1782 
.08987 

DENSITIES  AND  MELTING  POINTS  OF  SOME  COMMON  ELEMENTS 


NAME            DENSITY 
Aluminium     .     .     2.65 
Antimony      .     .     6.52 

MELTING 
POINT 

658.7° 
630.0° 

NAME 
Magnesium    . 
Manganese     . 

DENSITY 
.     1.74 
.     7.39 

MELTING 
POINT 

651.0° 
1260.0° 

Arsenic      .     .     . 

5.73 

Mercury    .     . 

.  13.56 

-38.7° 

Bismuth    .     .     . 

9.80 

271.0° 

Nickel  .     . 

.     8.9 

1452.0° 

Calcium     .     .     . 

1.55 

810.0° 

Phosphorus    . 

.     1.83 

44.0° 

Carbon,  diamond 

3.52 

Platinum  . 

.  21.50 

1755.0° 

Carbon,  graphite 

2.30 

<  4000.0° 

Potassium 

.     0.862 

62.3° 

Chromium      .     . 

6.50 

1520.0° 

Silicon  .     . 

.     2.3 

1420.0° 

Cobalt       .     .     . 

8.6 

1478.0° 

Silver   .     .     . 

.  10.5 

960.5° 

Copper      .     .     . 

8.93 

1083.0° 

Sodium 

.     0.97 

97.5° 

Gold     .... 

19.32 

1032.0° 

Sulfur  .     .     . 

.     2.06 

112.8° 

Iron      .... 

7.86 

1530.0° 

Tin  .... 

.     7.29 

231.9° 

Lead 

11.37 

327.0° 

Zinc      .     .     . 

.     7.10 

419.4° 

APPENDIX 


403 


SOLUBILITY  OF  VARIOUS  GASES  IN  WATER 


NAME  OF  GAS 

ATOLUME  ABSORBED  AT  0° 
UNDER  760  MM.   PRESSURE 
1  LITER  OF  WATER 

AND 
BY 

Ammonia    .           

1298.9  liters 

Hydrogen  clilorid6 

506.0  liters 

Sulfur  dioxide       

79.  79  liters 

Hydrogen  sulfide 

4  37  liters 

Carbon  dioxide     ... 

1.713  liters 

Oxygen 

0  0496  liters 

Hydrofren                        .          

0.0214  liters 

Nitrogen 

0  0233  liters 

TABLE  OF  SOLUBILITY  OF  VARIOUS  SOLIDS 


SUBSTANCE 

FORMULA 

WEIGHT  DISSOLVED  BY  100  cc.  OF  WATER  AT 

0° 

20° 

100° 

Calcium  chloride     . 

CaCl2 

59.5  g. 

74.5  g. 

159.0  g. 

Sodium  chloride 

NaCl 

35.70g. 

36.0  g. 

39.80  g. 

Potassium  nitrate    . 

KN03 

13.30  g. 

31.6  g. 

246.0  g. 

Copper  sulf  ate    .     . 

CuS04 

14.30g. 

21.  7g. 

75.4  g. 

Calcium  sulfate  .     . 

CaS04 

0.759g. 

0.203g. 

0.162  g. 

Calcium  hydroxide 

Ca(OH)2 

0.186  g. 

0.165  g. 

0.077  g. 

RELATION  OF  COMMON  UNITS  AND  METRIC  UNITS 


1  pound  (troy) 

1  pound  (avoirdupois) 

1  ounce  (avoirdupois) 

I  United  States  quart 

1  liter 

1  meter 

1  centimeter 

1  kilogram 


373.24  grams 

453.59  grams 

28.35  grams 

0.946  liters 

1.056  United  States,  quarts 

39.37  inches 

nearly  |  inch 

nearly  2^  pounds  avoirdupois 


INDEX 


Absolute  scale  of  temperature,  35 

Absolute  zero,  35 

Acetic  acid,  235 

Acetylene,  207 

Acid  anhydrides,  134 

Acids,  108  ;  binary,  113  ;  charac- 
teristics of,  107 ;  definition  of, 
108  ;  dibasic,  156  ;  familiar,  107  ; 
ionization  of,  108  ;  monobasic, 
156 ;  namingof,  113;  organic,  235; 
preparation  of,  138  ;  strength  of, 
111 ;  ternary,  113  ;  undissociated, 
108 

Affinity,  chemical,  9 

Agate,  262 

Agent,  bleaching,  150,  174;  cata- 
lytic, 152  ;  dehydrating,  155  ; 
oxidizing,  31  ;  reducing,  31 

Air,  86  ;  analysis  of,  88  ;  carbon 
dioxide  in,  88  ;  composition  of, 
86;  constancy  of  composition 
of,  91 ;  constituents  of,  essen- 
tial to  life,  87 ;  impure,  91  ; 
liquid,  92 ;  a  mixture,  90 ;  ni- 
trogen in,  88  ;  oxygen  in,  88 ; 
water  vapor  in,  87 

Air  saltpeter,  129 

Alchemists,  12 

Alchemy,  12 

Alcohol,  228  ;  absolute,  230 ;  de- 
natured, 230  ;  ethyl,  228  ;  grain, 
228  ;  methyl,  228  ;  wood,  228 

Alcoholic  liquors,  230 

Alcohols,  228 

Alizarin,  331 

Alkali,  107 

Alkali  metals,  276 

Alloys,  258  ;  steel,  348 

Aluminium,  322  ;  silicates  of,  333  ; 
hydrolysis  of  salts  of,  329  ;  in 
steel,  348 


Aluminium  bronze,  325 

Aluminium  hydroxide,  327  - 

Aluminium  oxide,  326 

Alums,  328 

Alundum,  326 

Amalgams,  362 

Amethyst,  202  ;  oriental,  326 

Ammonia,  121  ;  chemical  conduct 
of,  124  ;  decomposition  of,  135  ; 
preparation  of,  122  ;  properties 
of,  123  ;  volume  composition  of, 
127 

Ammoniacal  liquor,  210 

Ammonium,  122,  126 ;  compounds 
of,  288 

Ammonium  carbonate,  289 

Ammonium  chloride,  288 

Ammonium  hydroxide,  126 

Ammonium  radical,  126 

Ammonium  salts,  288 

Ammonium  sulfate,  289 

Ammonium  sulfides,  289 

Anesthetic,  132,  230 

Anhydride,  134 ;  carbonic,  198  ; 
phosphoric,  251  ;  silicic,  263 ; 
sulfuric,  148  ;  sulfurous,  148 

Anhydrite,  297 

Aniline,  232 

Aniline  dyes,  232 

Anions,  103 

Anode,  101 

Anthracene,  232 

Antimony,  255  ;  action  with  chlo- 
rine, 171;  alloys  of,  259;  com- 
pounds of,  256 

Apatite,  247,  306 

Aqua  ammonia,  122 

Aqua  regia,  178 

Aqueous  tension,  36 ;  table  of, 
402 

Argol,  237 


405 


406 


FIEST  COUESE  IN  CHEMISTRY 


Argon,  84 

Arrhenius,  Svante,  101 ;  portrait 
of,  102 

Arsenic,  252 ;  action  of  chlorine 
on,  171  ;  Marsh's  test  for,  254  ; 
molecular  weight  of,  189  ;  white, 
254 

Arsenic  acid,  255 

Arsenic  insecticide,  255 

Arsenic  snlfide,  255 

Arsenious  acid,  255 

Arsenious  oxide,  254 

Arsenious  sulfide,  255 

Arsenopyrite,  252 

Arsine,  253 

Artificial  silk,  225 

Asbestos,  316  ;  platinized,  392 

Atmosphere,  86 

Atom,  67  ;  definition  of,  69  ;  size 
of,  70 

Atomic  theory,  67 

Atomic  weights,  70  ;  accurate  de- 
termination of,  191 ;  from  com- 
bining weights,  189  ;  relation  of, 
to  properties  of  elements,  162  ; 
steps  in  determining,  191 

Avogadro's  hypothesis,  185 ;  and 
molecular  weights,  186 

Azote,  80 

Babbitt  metal,  259 

Bakelite,  233 

Baking,  329 

Baking  powders,  330 

Barite,  307 

Barium,  297,  307 

Barium  chloride,  307 

Barium  oxide,  307 

Barium  peroxide,  307 

Barium  sulfate,  307 

Bases,  109  ;  characteristics  of,  108  ; 

definition  of,  109  ;  familiar,  108  ; 

ionization  of,  109  ;    naming  of, 

113  ;  strength  of,  111 
Basic  salts,  257 
Bauxite,  322 
Beer,  230 
Benzene,  232 
Benzine,  205 
Benzoic  acid,  232 
Berzelius,  48 


Bessemer  converter,  344 

Bessemer  process,  344 

Birkeland  and  Eyde  process,  128 

Bismuth,  256  ;  alloys  of,  259  ;  hy- 
drolysis of  salts  of,  257 

Bismuth  chloride,  257 

Bismuth  nitrate,  257 

Bismuth  oxide,  257 

Bismuth  subnitrate,  257 

Bismuthyl  chloride,  257 

Blast  furnace,  341 

Bleach,  303 

Bleaching,  174;  agents,  174;  by 
chlorine,  173  ;  by  hydrogen  per- 
oxide, 52  ;  by  sulfurous  ac\d, 
150 

Bleaching  powder,  171,  303 

Blowpipe,  oxyacetylene,  209  ;  oxy- 
hydrogen,  31 

Blue  printing,  352 

Bluestone,  361 

Boiler  scale,  315 

Boiling  point,  56 

Bone  ash,  248 

Bone  black,  197 

Borax,  270 

Bordeaux  mixture,  361 

Boric  acid,  269 

Bornite,  357 

Boron,  269  ;  acids  of,  269 

Boyle,  Robert,  33 

Brass,  360 

Brazing,  270,  370 

Bread  making,  231 

Brick,  333  ;  vitrified,  334 

Brimstone,  142 

Bromides,  181 

Bromine,  178 

Bronze,  360  ;  aluminium,  360 

Bunsen,  Robert  (portrait),  277 

Burners,  Bunsen,  217  ;  gas-stove, 
217 

Burning,  1  ;  increase  in  weight 
during,  2  ;  of  iron,  3  ;  of  phos- 
phorus, 3 

Butane,  203 

Butter,  238 

Butter  fat,  238 

Butyric  acid,  235 

Butyrin,  238 

By-product,  18,  282 


INDEX 


407 


Cadmium,  313,  321 

Caesium,  276 

Caleite,  298 

Calcium,  297 

Calcium  acid  carbonate,  299 

Calcium  acid  sulfite,  306 

Calcium  carbide,  198,  207,  305 

Calcium  carbonate,  298 

Calcium  cyanamide,  305 

Calcium  fluoride,  306 

Calcium  hydroxide,  302  ;  action  of 
carbon  dioxide  on,  201 

Calcium  oxide,  301 

Calcium  phosphate,  306 

Calcium  sulfate,  304 

Calcium  sulfide,  306 

Calomel,  363 

Calorie,  61 

Calorimeter,  61 ;  respiration,  245 

Candles,  294 

Caramel,  222 

Carbides,  198 

Carbohydrates,  220 

Carbolic  acid,  232 

Carbon,  193  ;  amorphous,  195  ; 
chemical  conduct  of,  198  ;  prop- 
erties of,  197 

Carbon  retort,  211 

Carbon  dioxide,  198 ;  action  of, 
on  calcium  hydroxide,  201  ;  and 
plant  life,  89  ;  chemical  conduct, 
199  ;  in  air,  87  ;  variation  of,  in 
air,  89 

Carbon  disulfide,  156 

Carbon  monoxide,  202 

Carbon  tetrachloride,  207 

Carbona,  207 

Carbonates,  201 

Carbonic  acid,  200  ;  salts  of,  201 

Carborundum,  260 

Carnallite,  286 

Carnotite.  379 

Casein,  222 

Cassiterite,  369 

Catalysis,  152  ;  examples  of,  152 

Catalytic  agent,  152,  392 

Catalyzers,  152  ;  negative,  153 

Cathode,  101 

Cations,  103 

Caustic  potash,  286 

Caustic  soda,  278 


Cavendish,  24,  40 

Caves,  formation  of,  299 

Celluloid,  225 

Cellulose,  220,  225 

Cement,  335 

Cementite,  339 

Cerium,  396 

Cerium  oxide,  213 

Chalcedony,  262 

Chalcocite,  357 

Chalcopyrite,  357 

Chalk,  299;  French,  317 

Chalybeate  water,  350 

Chamberlain-Pasteur  filter,  43 

Charcoal,  196  ;  animal,  197 

Charles,  law  of,  35 

Cheese,  222 

Chemical  action,  9,  61 

Chemical  affinity,  9 

Chemical  change,  9 

Chemical  energy,  61 

Chile  saltpeter,  181 

Chlorauric  acid,  395 

Chloric  acid,  178 

Chlorides,  171,  178 

Chlorine,  169  ;  action  of,  on  bases, 
287 ;  action  of,  as  a  disinfectant, 
174 ;  action  of,  on  elements,  171  ; 
action  of,  on  hydrogen,  172  ;  ac- 
tion of,  on  water,  172  ;  bleaching 
action  of,  173;  compounds  of, 
with  oxygen  and  hydrogen,  178 

Chlorine  family,  166 

Chloroform,  207 

Chlorophyll,  313,  338 

Chloroplatinic  acid,  391,  393 

Chlorous  acid,  178 

Chromates,  388 

Chromic  acid,  388 

Chromic  hydroxide,  387 

Chromic  oxide,  387 

Chromic  salts,  387 

Chromite,  387 

Chromium,  387 

Cinnabar,  362,  363 

Citric  acid,  237 

Clay,  333 

Clay  products,  333 

CoaX  196 

Coal  gas,  210 

Coal  tar,  210  ;  derivatives  of,  232 


408 


FIRST  COURSE  IN  CHEMISTRY 


Coal-tar  compounds  in  food,  233 

Cobalt,  354  ;  compounds  of,  356 

Cobalt  oxide,  355 

Coinage,  395 

Coke,  196 

Cold  storage,  125,  231 

Colemanite,  270 

Collodion,  225 

Colloidal  solutions,  99 

Colloidal  suspensions,  99 

Colloids,  99 

Combining  weights,  65 ;  atomic 
weights  from,  189  ;  law  of,  66 ; 
standard  for,  66 

Combustion,  20 ;  effect  of,  on  com- 
position of  air,  89  ;  in  oxygen, 
93 ;  spontaneous,  21 

Compound,  definition  of,  8  ;  num- 
ber of,  12 

Concrete,  337  ;  reinforced,  337 

Condensite,  233 

Copper,  357 ;  action  of  chlorine 
on,  171  ;  alloys  of,  360  ;  blister, 
359  ;  compounds  of,  361;  con- 
verter, 358  ;  refining  of,  359  ; 
ruby,  361 

Copper  oxide,  reduction  of,  by  hy- 
drogen, 31 

Copperas,  349 

Corn  sirup,  223 

Corrosive  sublimate,  363 

Corundum,  326 

Cotton  fiber,  226 

Coumarin,  233 

Cream  of  tartar,  237 

Cresol,  232 

Crisco,  239 

Cryolite,  166,  322 

Crystallography,  58 

Crystals,  58 

Cupric  compounds,  361 

Cupric  oxide,  361 

Cupric  sulfate,  361 

Cupric  sulfide,  362 

Cuprite,  357,  361 

Cuprous  compounds,  361 

Cuprous  oxide,  361 

Curie,  Madame,  381 

Cyanide  process  for  gold,  394 

Cyanides,  203 

Cyanogen,  203 


Dalton,  John,  67 

Davy,  Sir  Humphry,  169  ;  portrait 
of,  206 

Decay,  89 

Definite  composition,  law  of,  51 

Dehydrating  agent,  155 

Densities  of  elements,  table  of, 
402 

Developing,  367 

Dewar  flask,  93 

Dextrin,  223 

Dextrose,  220,  222,  223 

Diamond,  193  ;  artificial  prepara- 
tion of,  194;  Cullinan,  193; 
Kohinoor,  194 

Dibasic  acids,  156 

Dichromic  acid,  388 

Dietary  standards,  243 

Diets,  calculation  of,  244 

Displacement  series,  112 

Distillation,  41 

Dolomite,  315 

Dore"  bars,  373 

Dry  cleaning,  205 

Dryer,  376 

Dumas,  48 

Dust  explosions,  94 

Dyeing,  330 

Dyes,  330  ;  aniline,  232,  331 

Dynamite,  296  ;  gelatin,  296 

Earth's  crust,  composition  of,  9 

Ebonite,  145 

Effervescence,  200 

Eggs,  preservation  of,  265 

Electric  furnaces,  215 

Electrochemical  industries,  274 

Electrodes,  101 

Electrolysis,  101  ;  and  ionization, 
104 ;  of  sodium  chloride,  104 ; 
of  water,  105 

Electrolytes,  101 

Electrons,  382 

Electroplating,  365;  with  nickel, 
355  ;  with  silver,  365 

Electroscope,  380 

Electrotyping,  360 

Elements,  8  ;  classification  of ,  159  ; 
essential  to  life,  10  ;  families  of, 
163  ;  in  human  body,  10 ;  mo- 
lecular weight  of,  189;  names 


INDEX 


409 


of,  10 ;  number  of,  9 ;  occur- 
rence of,  10  ;  symbols  of,  11 

Emery,  326 

Emulsions,  99 

Enamels,  269 

Energy,  60  ;  chemical,  61  ;  conser- 
vation of,  60 ;  transformation  of, 
60 

Epsom  salt,  316 

Equations,  74 ;  and  calculations, 
77  ;  and  volumes  of  gases,  191 

Equilibrium,  135 ;  in  solution, 
136 

Etching,  169 

Ethane,  203 

Ether,  230 

Eudiometer,  47,  50 

Evaporation,  56 

Explosions  due  to  dust,  94 

Explosives,  294 

Families  in  periodic  groups,  163 

Faraday,  57  ;  portrait  of,  58  ;  and 
liquefaction  of  gases,  57 

Fats,  237 

Fatty-acid  series,  235 

Feldspar,  264 

Fermentation,  acetic,  236 ;  alco- 
holic, 228,  329  ;  lactic,  222 

Ferric  compounds,  349 ;  reduction 
of,  351 

Ferric  chloride,  350 

Ferric  hydroxide,  350 

Ferric  nitrate,  353 

Ferric  sulfate,  353 

Ferrochromium,  387 

Ferromanganese,  385 

Ferrotitanium,  397 

Ferrous  carbonate,  350 

Ferrous  compounds,  349 ;  oxida- 
tion of,  351 

Ferrous  sulfate,  349 

Ferrous  sulfide,  349 

Ferrovanadium,  397 

Fertilizers,  309  ;  commercial,  310  ; 
nitrogen  in,  309 ;  potassium  in, 
310 

Fire  curtains,  259 

Fire  damp,  206 

Fire  extinguishers,  200 

Flame  reactions,  289 


Flames,  215;  candle,  217;  condi- 
tions necessary  for,  215;  struc- 
ture of,  216 

Flash  lights,  314 

Flint,  262 

Fluorapatite,  166 

Fluorides,  168 

Fluorine,  166 

Fluorite,  166,  297,  306 

Flux,  340 

Foods,  240 ;  bleaching  of,  174 ; 
composition  of,  240 ;  cost  of  and 
nutritive  value,  245 ;  energy 
value  of,  242 ;  fuel  value  of, 
242  ;  function  of,  240  ;  necessary 
for  health,  243  ;  table  of,  241 

Formaldehyde,  228 

Formalin,  228 

Formic  acid,  235  ;  preparation  of 
carbon  monoxide  from,  202 

Formula  weights,  73 

Formulas,  72  ;  facts  expressed  by, 
72 

Franklinite,  317 

Freezing  point,  56 

Fruit  trees,  spraying  of,  144 

Fruits,  bleaching  of,  150 

Fuel  gases,  composition  of,  213 

Fuels,  210 ;  ""-products  of  combus- 
tion of,  214 

Fuller's  earth,  333 

Furnaces,  electric  :  arc,  215 ;  re- 
sistance, 215 

Galenite,  375 

Galvanized  iron,  319 

Gas,  coal,  210 ;  natural,  212  ;  pro- 
ducer, 212  ;  water,  211 

Gas  laws,  33 

Gas  mantles,  213 

Gases,  collection  of,  17  ;  liquefac- 
tion of,  57  ;  table  of  weights  of, 
402  ;  volumes  of,  from  equations, 
191 ;  weight  of  a  liter  of,  192 

Gasoline,  205 

Gay-Lussac,  Joseph,  35,  36 

Gems,  artificial,  326 

Glacial  acetic  acid,  235 

Glass,  265;  blowing  of,  265;  col- 
oring of,  267  ;  etching  of,  169 ; 
fluorescent,  379;  Jena,  267; 


410 


FIKST  COURSE  IN  CHEMISTRY 


milky,  268;  molding  of,  205; 
nature  of,  268:  plate,  266;  se- 
lenium in,  398  ;  varieties  of,  267 ; 
window,  266 

Glauber's  salt,  280 

Glucose,  223 

Glycerin,  238,  294 

Gold,  393;  fool's,  350;  recovered 
from  copper,  359 

Gold  coin,  360 

Goldschmidt  reduction  process,  325 

Gram-atomic  weight,  74 

Gram-molecular  volume  of  gases, 
188 

Gram-molecular  weights.  73 

Granite,  233,  264 

Graphite,  195 

Grease  spots,  removal  of,  353 

Gun  metal,  360 

Guncotton,  225 

Gunpowder,  294  ;  smokeless,  295 

Gypsum,  304 

Haber  process  for  manufacture  of 
ammonia,  123 

Haemoglobin,  338 

Hall,  Charles  Martin,  323;  por- 
trait of,  322 

Halogens,  166 

Hard  water,  293,  300 

Health  and  ventilation,  91 

Heat,  59  ;  a  form  of  energy,  60  ; 
measurement  of,  61 ;  of  forma- 
tion, 46  ;  of  fusion,  56 ;  of  oxi- 
dation and  combustion,  21  ;  of 
reaction,  representation  of,  76 ; 
of  solidification,  56 ;  transfor- 
mations of,  59  ;  unit  of,  61 

Helium,  84 

Hematite,  339 

Henry's  law,  99 

Heptane,  203 

Hexane,  203 

Human  body,  composition  of,  10 

Hydrates,  46,  271 

Hydraulic  mining,  394 

Hydrides,  28 

Hydriodic  acid,  182 

Hydrobromic  acid,  181 

Hydrocarbons,  203 

Hydrochloric  acid,  177;  salts  of,  178 


Hydrocyanic  acid,  203 

Hydrofluoric  acid,  168 ;  etching 
with,  169 

Hydrogen,  24 ;  action  of,  on  chlo- 
rine, 172  ;  burning  of,  29  ;  chem- 
ical conduct  of,  28  ;  preparation 
of,  24 ;  preparation  of,  from 
acids,  26  ;  properties  of,  28  ;  uses 
of,  31 

Hydrogen  bromide,  180 

Hydrogen  chloride,  175  ;  action  of, 
as  an  acid,  177  ;  action  of,  on 
oxidizing  agents,  177  ;  composi- 
tion of,  176 

Hydrogen  fluoride,  168  ;  electrol- 
ysis of,  166 

Hydrogen  iodide,  182 

Hydrogen  peroxide,  51 

Hydrogen  sulfate,  154 

Hydrogen  sulfide,  145 

Hydrolysis,  138 

Hydrosulf uric  acid,  146 ;  salts  of, 
147 

Hydroxyl  radical,  112 

Hydroxyl  ion,  109 

Hypochlorous  acid,  178 

Hypothesis,  Avogadro's,  185 

Ice,  manufacture  of,  124 

Iceland  spar,  299 

Ilmenite,  397 

Indicators,  107 

Indigo,  331 

Industries,  electrochemical,  274 

Infusorial  earth,  262 

Ink  stains,  353 

Inks,  353 

Insecticides,  arsenic,  255 

Iodides,  183 

Iodine,  181  ;  tincture  of,  182 

lodoform,  182,  207 

lonization,  and  boiling  point,  104  ; 
and  electrolysis,  104  ;  extent  of, 
110;  and  freezing  point,  104; 
theory  of,  101 

Ions,  definition  of,  102;  charges 
on  and  valence,  119  ;  electrical 
charge  of,  102 ;  formation  of, 
102  ;  and  properties  of  electro- 
lytes, 106 

Iridiurn,  398 


INDEX 


411 


Iron,  338  ;  action  of,  on  steam,  26  ; 
cast,  341 ;  compounds  of,  341);  gal- 
vanized, 318  ;  pig,  342  ;  wrought, 
343 

Iron  alum,  353 

Iron  disulphide,  350 

Iron  ore,  339 

Isomeric  compounds,  220 

Jasper,  262 

Jewels,  artificial,  267 

Kainite,  286,  310 

Kaolin,  264,  333 

Kelp,  181 ;  potassium  in,  285 

Kerosene,  205 

Kieserite,  286 

Kindling  temperature,  21 

Kinetic  molecular  theory,  37 

Krypton,  84 

Lactic  acid,  222 

Lactose,  220,  222 

Lakes,  332,  377 

Lampblack,  197 

Laughing  gas,  132 

Lavoisier,  5,  55  ;  portrait  of,  Fron- 
tispiece 

Law,  67  ;  of  Boyle,  33  ;  of  Charles, 
35 ;  of  combining  weights,  66 ; 
of  conservation  of  energy,  60 ; 
of  conservation  of  matter,  55 ;  of 
definite  composition,  51  ;  of  Gay- 
Lussac,  35 ;  of  Henry,  99  ;  of  mul- 
tiple proportion,  52 ;  natural,  67 ; 
periodic,162  ;  of  volumes,  183, 185 

Lead,  371  ;  desilvering  of,  373 ; 
hard,  372  ;  red,  374  ;  silver-bear- 
ing, 372  ;  soft,  372  ;  softening  of, 
372  ;  sugar  of,  235  ;  white,  375 

Lead  arsenate,  255 

Lead  carbonate,  375 

Lead  oxides,  374 

Lead  peroxide,  375 

Lead  sulfide,  375 

Leather,  388 

Leblanc,  Nicolas,  282 

Leblanc  process,  281 

Levulose,  220,  222 

Liebig,  Justus,  portrait  of,  45 

Lime,  301 ;  air-slaked,  302  ;  chloride 


of,  303 ;    commercial  production 

of,  302  ;  slaked,  302 
Limekiln,  302 
Limelight,  32,  301 
Lime-nitrogen,  306 
Limestone,  298  ;  dolomitic,  315 
Lime-sulfur  spray,  144 
Limewater,  302 
Liming  soils,  110 
Limonite,  339 
Linseed  oil,  376 
Liquefaction  of  gases,  57 
Liquid-air  machines,  58 
Litharge,  374 
Lithium,  276 
Lithopone,  320 
Litmus,   action  of  acids  on,   107  ; 

action  of  bases  on,  108 
Lockyer,  84 
Lunar  caustic,  366 
Lye,  279 

Magnalium,  325 

Magnesia,  314 

Magnesite,  315 

Magnesium,  313 ;  hydrolysis  of 
salts  of,  315 

Magnesium  carbonate,  315 

Magnesium  hydroxide,  314 

Magnesium  oxide,  314 

Magnesium  silicates,  316 

Magnesium  sulfate,  316 

Magnetite,  339 

Malachite,  357 

Malt,  229 

Maltose,  220 

Manganese,  385 

Manganese  dioxide,  385  ;  action  of 
hydrochloric  acid  on,  170 

Manganous  salts,  386 

Marble,  299 

Marsh  gas,  206 

Marsh's  test  for  arsenic,  254 

Mass  action,  136 

Matches,  249 

Matte,  358 

Matte  furnace,  358 

Matter,  55 ;  amorphous,  58 ;  con- 
servation of,  55  ;  crystalline,  58  ; 
states  of,  55 

Meerschaum,  316 


412 


FIRST  COURSE  IN  CHEMISTRY 


Melaconite,  357 

Melting  point,  56 

Melting  point  of  elements,  table  of, 
402 

Mendele"eff,  159  ;  portrait  of,  160 

Mercerized  cotton,  225 

Mercuric  chloride,  363 

Mercuric  compounds,  363 

Mercuric  oxide,  decomposition  of, 
by  heat,  4,135;  preparationof, 363 

Mercuric  sulfide,  363 

Mercurous  chloride,  363 

Mercurous  compounds,  363 

Mercury,  362 

Metaboric  acid,  269 

Metallurgy,  272 

Metals,  the,  272 ;  compounds  of,  273 

Metaphosphoric  acid,  252 

Metasilicic  acid,  264 

Metastannic  acid,  370 

Meteorites,  338 

Methane,  203,  206  ;  halogen  deriv- 
atives of,  207 

Methods,  laboratory  and  commer- 
cial, 18 

Mica,  264,  333 

Microcosmic  salt,  252 

Milk,  100;  composition  of,  222; 
souring  of,  222 

Minium,  374 

Mixed  salts,  252 

Moissan,  Henri,  166,  167,  194 

Molasses,  221 

Molecular  weights,  185  ;  and  Avo- 
gadro's  hypothesis,  186  ;  of  ele- 
ments, 189  ;  oxygen  a  standard 
for,  186;  from  weight  of  1  liter,  188 

Molecules,  37  ;  definition  of,  69 

Molybdenum,  397 

Monazite  sand,  213 

Monobasic  acids,  156 

Mordants,  332 

Morley,  49 

Mortar,  303 

Moth  balls,  232 

Multiple  proportion,  law  of,  52 

Muriatic  acid,  176 

Naphtha,  205 
Naphthalene,  232 
Nascent  state,  175 


Natural  gas,  206,  212 

Neon,  84 

Neutralization,  109,  138;  illustra- 
tion of,  110 

Nickel,  354 ;  as  catalytic  agent, 
239  ;  compounds  of,  356  ;  recov- 
ered from  copper,  359 

Nickel  coin,  360 

Nickel  plating,  355 

Nitrates,  131 

Nitric  acid,  127 ;  acid  action  of, 
129  ;  action  of,  on  metals,  130  ; 
commercial  preparation  of,  128  ; 
decomposition  of,  on  heating,  129; 
fuming,  130  ;  oxidizing  action  of, 
130 ;  preparation  of,  from  air, 
128  ;  salts  of,  131 ;  uses  of,  130 

Nitric  oxide,  132 

Nitrides,  83 

Nitrites,  131 

Nitrobenzene,  232 

Nitrocellulose,  225,  295 

Nitrogen,  80 ;  acids  of,  127 ;  as- 
similation of,  by  plants,  83 ; 
chemical  conduct  of,  82  ;  oxides 
of,  131  ;  preparation  of,  from  air, 
81 ;  preparation  of,  from  ammo- 
nium nitrite,  82  ;  in  soils,  121  ; 
utilization  of  atmospheric,  311 

Nitrogen  dioxide,  133 ;  formed 
from  nitric  acid,  129 

Nitrogen  pentoxide,  131,  134 

Nitrogen  tetroxide,  133 

Nitrogen  trioxide,  131,  134 

Nitroglycerin,  294,  295 

Nitrous  acid,  131 

Nitrous  oxide,  132 

Normal  salts,  156 

Nutrition,  animal  and  plant,  246 

Oil  of  vitriol,  153 
Oil  wells,  204 

Oils,  237  ;  cracking  of,  206 ;  lubri- 
cating, 205 
Oleic  acid,  237 
Olein,  238 
Oleomargarine,  238 
Onyx,  262 
Opal,  262 

Open-hearth  process,  345 
Organic  acids,  235 


INDEX 


413 


Orpiment,  252 

Orthosilicic  acid,  264 

Oxalic  acid,  preparation  of  carbon 

monoxide  from,  202 
Oxidation,  20,  351  ;  of  ferrous  salts, 

351 

Oxides,  20 
Oxidizing  agent,  31 
Oxyacetylene  blowpipe,  209 
Oxygen,  14 ;  chemical  conduct  of, 
19  ;    commercial  preparation  of, 
17;  discovery  of ,  14 ;  importance 
of,  22  ;  occurrence  of,  14  ;  prep- 
aration of,  from  mercuric  oxide, 
15;  preparation  of,  from  potas- 
sium chlorate,  15 ;  properties  of, 
19 ;     standard     for     molecular 
weights,  186 

Oxygen  atom,  weight  of,  186 
Oxygen  molecule,  weight  of,  187 
Oxyhydrogen  blowpipe,  32 
Ozone,  62 ;  molecular  weight  of,  189 

Paints,  376 
Palladium,  398 
Palmitic  acid,  235 
Palmitin,  238 
Paper,  226 
Paraffin,  205 
Paris  green,  255 
Paste  jewels,  267 
Pearls,  298 
Pentane,  203 
Perchloric  acid,  178 
Periodic  grouping,  159 
Periodic  law,  162  ;  value  of,  164 
Periodic  table,  161 
Permanganate,  386 
Permanganic  acid,  386 
Petroleum,  204  ;  refining  of,  205 
Pewter,  370 
Phenol,  232 

Philosopher's  stone,  12 
Phlogiston,  6 
Phosphates,  252,  309 
Phosphine,  250 
Phosphonium  chloride,  251 
Phosphorescence,  249 
Phosphoric  acid,  251 
Phosphorite,  247,  297,  306 
Phosphorous  acid,  252 


Phosphorus,  247 ;  molecular  weight 
of,  189  ;  oxides  of,  251  ;  red,  249; 
white,  248 

Phosphorus  family,  247 

Phosphorus  pentachloride,  252 

Phosphorus  pentoxide,  251 

Phosphorus  sesquisulfide,  250,  252 

Phosphorus  trichloride,  252 

Photography,  366 

Pitchblende,  379 

Placer  mining,  394 

Plant  food,  309 

Plants,  effect  of,  on  composition  of 
air,  89  ;  silica  in,  260 

Plaster,  303 

Plaster  of  Paris,  304 

Platinum,  391 ;  catalytic  action  of ,  153 

Platinum  sponge,  391 

Pneumatic  trough,  17 

Polyboric  acids,  269 

Polysilicic  acids,  264 

Porcelain,  334 

Portland  cement,  335 

Potassium,  284  ;  in  sea  plants,  285 ; 
properties,  286 

Potassium  alum,  328 

Potassium  bicarbonate,  288 

Potassium  bisulfate,  288 

Potassium  bromide,  288 

Potassium  carbonate,  288 

Potassium  chlorate,  287 ;  decompo- 
sition of,  75 

Potassium  chloride,  288 

Potassium  chromate,  388 

Potassium  cyanide,  203 

Potassium  dichromate,  388 

Potassium  f erricyanide,  352 

Potassium  f  errocyanide,  351 

Potassium  hydroxide,  286 

Potassium  iodide,  183,  288 

Potassium  nitrate,  287 

Potassium  permanganate,  386 

Potassium  sulfate,^88 

Pottery,  334 

Precipitate,  138 

Precipitation,  137 

Preservatives,  231 

Pressure,  standard,  36 

Priestley,  Joseph,  4f  portrait  of,  15 

Producer,  gas,  212 

Propane,  203 


414 


FIRST  COURSE  IX  CHEMISTRY 


Properties,  18 
Protein,  121,  239 
Prussian  blue,  352 
Prussiate  of  potash,  red,  352  ;  yel- 
low, 351 

Prussic  acid,  203 
Puddling  furnace,  343 
Pyrite,  350 
Pyrolusite,  385 
Pyrophosphoric  acid,  252 

Quartz,  262 
Quartz  glass,  263 
Quartz  mining,  394 
Quicklime,  301 

Radical,  glyceryl,  238;  hydroxyl, 
112;  nitrate,  112;  sulfate,  112 

Radicals,  112 

Radioactivity,  379 

Radiograph,  380 

Radium,  discovery  of,  380;  disinte- 
gration of,  382  ;  energy  of,  383  ; 
and  medicine,  383  ;  origin  of, 
382 ;  quantity  available,  381 

Ramsay,  Sir  William,  84  ;  portrait 
of,  83 

Rare  earths,  396 

Rayleigh,  84 

Reactions,  75 ;  completed,  136 ; 
heat  of,  76  ;  reversible,  135 

Realgar,  252 

Reducing  agent,  31 

Reduction,  30 ;  of  ferric  salts,  351 ; 
Goldschmidt  process  of,  325 

Reference  books,  400 

Rennin,  222 

Replacing  power  of  atoms,  118 

Respiration,  effect  of,  on  compo- 
sition of  air,  89 

Reversible  reactions,  6,  135 

Roasting,  317 

Rock  phosphate,  306 

Rocks,  weathering  of,  by  carbon 
dioxide,  90 

Rubber,  vulcanizing  of,  145 

Rubidium,  276 

Ruby,  326 

Rust  stains,  removal  of,  354 

Rutherford,  80 

Rutile,  397 


Saccharine,  233 

Safety  lamp,  206 

Sal  ammoniac,  288 

Sal  soda,  282 

Salt,  279;  rock,  279 

Saltpeter,  287 ;  air,  129 ;  Chile,  283 

Salts,  109;  acid,  156;  basic,  258; 
characteristics  of,  110;  mixed, 
252  ;  naming  of,  114;  normal,  156; 
preparation  of,  273  ;  solubilities 
of,  274;  Stassfurt,  284,  286 

Sand,  262 

Sandstone,  262 

Saponification,  292 

Sapphire,  326 

Scheele,  Karl,  14,  80,  169 

Scheele's  green,  255 

Scouring  soap,  262 

Selenium,  398 

Serpentine,  316 

Sewage-disposal  plant,  22 

Shot,  253 

Siderite,  339 

Silica,  262 ;  action  of  hydrofluoric 
acid  on,  168,  263 

Silicates,  263,  264 

Silicic  acids,  264 

Silicides,  261 

Silicon,  260 

Silicon  dioxide,  262 

Silk,  fiber,  226;  artificial,  225; 
stains  on,  354 

Silver,  364;  German,  360;  recovered 
from  copper,  359 

Silver  bromide,  180 

Silver  chloride,  366 ;  precipitation 
of,  138 

Silver  coin,  360 

Silver  nitrate,  366 

Silver  plating,  365 

Silver  sulfide,  366 

Slag,  340 

Smalt,  356 

Smithsonite,  317 

Smoke  consumer,  218 

Smoke  prevention,  217 

Smokeless  powder,  295 

Soap,  composition  of,  291 ;  manu- 
facture of,  292;  properties  of,  293; 
scouring,  262  ;  varieties  of,  293 

Soap  powders,  293 


INDEX 


415 


Soapstone,  316 

Soda,  282 ;  baking,  283 ;  bicarbonate 
of,  283 ;  caustic,  278 ;  washing,  282 

Soda  ash,  282 

Soda  water,  200 

Sodium,  270 

Sodium  benzoate,  232,  233 

Sodium  carbonate,  280 

Sodium  chloride,  279 ;  electrolysis 
-  of,  104,  171 

Sodium  cyanide,  203 

Sodium  family,  276 

Sodium  hydrogen  carbonate,  283 

Sodium  hydroxide,  278 

Sodium  hyposulfite,  284 

Sodium  nitrate,  121,  283 

Sodium  peroxide,  278 

Sodium  phosphate,  252,  284 

Sodium  sulfate,  280 

Sodium  sulfite,  284 

Sodium  thiosulfate,  284 

Softening  of  water,  300 

Soils,  liming  of,  311 

Solder,  hard,  370;  soft,  370 

Soldering,  370 

Solubility,  96 ;  conditions  affecting, 
97;  effect  of  pressure  on,  99; 
effect  of  temperature  on,  98; 
table  of,  98 

Solubility  of  gases,  table  of,  403 

Solubility  of  salts,  274 

Solubility  of  solids,  table  of,  403 

Solute,  96 

Solutions,  96 ;  boiling  point  of, 
100;  classes  of,  97;  colloidal, 
99;  completion  of  reactions  in, 
137  ;  conductivity  of,  103 ;  elec- 
trolysis of,  101 ;  equilibrium  in, 
136 ;  freezing  point  of,  101 ;  prop- 
erties of,  100  ;  saturated,  96  ; 
supersaturated,  97 

Solvay  process,  281 

Solvent,  96  ;  effect  of,  on  solubility, 
98 

Spelter,  317 

Sphalerite,  317,  320 

Spontaneous  combustion,  21 

Spray,  lime-sulfur,  144 

Stains,  removal  of,  353 

Stalactites,  300 

Stalagmites,  300 


Standard  conditions,  36 

Stannic  chloride,  371 

Stannic  compounds,  371 

Stannous  chloride,  371 

Starch,  220,  223  ;  manufacture  of, 
224  ;  varieties  of,  225 

Stassfurt  salts,  179,  284,  286 

Stearic  acid,  235 

Stearin,  238 

Steel,  344;  hardening  of,  347;  prop- 
erties of,  347  ;  tempering  of,  347 

Steel  alloys,  348 

Steel  purifiers,  348 

Stibine,  256 

Stibnite,  255 

Strontium,  297,  307 

Sucrose,  220,  221 

Sugar,  221;  grape,  223;  invert, 
222 ;  milk,  222 

Sugar  of  lead,  235 

Sulfates,  156 

Sulfides,  146,  147  ;  use  of,  in  analy- 
sis, 147 

Sulfites,  151 

Sulfur,  140 ;  burning  of,  in  oxygen, 
19 ;  chemical  conduct  of,  143 ; 
compounds  of,  with  hydrogen, 
145;  extractionof,  140;  flowersof, 
142 ;  occurrence  of,  140 ;  oxides 
of,  148 ;  properties  of,  143 ;  roll, 
142 ;  uses  of,  144  ;  varieties  of ,  142 

Sulfur  dioxide,  148 

Sulfur  trioxide,  151 

Sulf  uric  acid,  153 ;  action  of, on  met- 
als, 155 ;  action  of,  on  salts,  155 ; 
action  of,  on  water,  155;  manu- 
facture of,  153 ;  oxidizing  prop- 
erties of,  154;  salts  of,  156 

Sulfurous  acid,  149 ;  bleaching  by, 
150;  salts  of,  151 

Supplementary  books,  400 

Sylvite,  286 

Symbol  weight,  74 

Symbols,  11 

Talc,  316 
Tannic  acid,  353 
Tanning,  389 
Tartari'c  acid,  237 
Tellurium,  396 
Temperature,  standard,  36 


416 


FIKST  COUESE  IN  CHEMISTRY 


Tetraboric  acid,  270 

Textile  fibers,  226 

Theme  writing  in  chemistry,  400 

Theories,  67 

Theory,  atomic,  67 ;  value  of,  69 

Theory  of  ionization,  101 

Thermite,  326 

Thermite  welding,  325 

Thermos  bottle,  93 

Thorium,  radioactive,  384 

Thorium  oxide,  213,  397 

Tin,  369 ;  block,  369 ;  compounds 
of,  371 ;  crystals,  371 ;  oxy mu- 
riate of,  371 

Tin  foil,  369 

Tin  plate  370 

Tinstone,  269 

Titanium,  397 

Toluene,  232 

Topaz,  326 

Trinitrotoluene,  296 

Tungsten,  397 

Type  metal,  259 

Undercooled  liquids,  56 

Units,    relation    of    common    and 

metric,  403 

Uranium,  379 ;  radioactivity  of,  379 
Uranyl  nitrate,  379 
Urea,  246 

Valence,  116;  and  charge  on  ions, 
119;  classification  according  to, 
116;  definition  of,  116;  deter- 
mination of,  117;  table  of,  119; 
variable,  118 

Vanadium,  397 

Vanilla  extracts,  233 

Vanillin,  233 

Vaporization,  56 

Vaseline,  204 

Ventilation,  91,  214 

Vermilion,  364 

Vinegar,  236;  mother  of,  236; 
varieties  of,  237 

Vitriol,  blue,  361 ;  green,  349 ;  oil 
of,  153,  154 ;  white,  320 

Volume  of  gases,  and  pressure,  33 

Volume  of  gases,  and  temperature, 
34 

Volumes,  law  of,  183 


Vulcanite,  145 
Vulcanized  rubber,  145 

Water,  action  of, on  iron,  26 ;  chem- 
ical conduct  of,  46  ;  city  filtration 
of,  44;  composition  of,  by  vol- 
ume, 47  ;  decomposition  of,  by 
metals,  24 ;  determination  of  ex- 
act composition  of,  46  ;  distilla- 
tion of,  41 ;  electrolysis  of,  17, 
105  ;  exact  composition  of,  50  ; 
hard,  41 ;  and  health,  41,  45 ;  min- 
eral matter  in,  41  ;  organic  matter 
in,  41 ;  properties  of,  45 ;  purifi- 
cation of,  41,  327  ;  purification 
of,  by  boiling,  42  ;  purification  of, 
by  distillation,  41 ;  purification 
of,  by  filtration,  43  ;  self-purifica- 
tion of,  45 ;  soft,  41 ;  softening 
of,  300 ;  volume  of  vapor  of,  com- 
pared with  volume  of  hydrogen 
and  oxygen,  50 

Water  gas,  211 ;  enriched,  212 

Water  glass,  265 

Water  of  hydration,  271 

Water  vapor  in  air,  87 

Weathering  of  rocks,  90 

Weight  of  gases,  table  of,  403 

Weight  of  1  liter  of  a  gas,  192 

Weights,  atomic,  70 

Welding,  thermite  process  for,  325 

Whey,  222 

Whisky,  230 

Wine,  230 

Wollastonite,  297 

Wood,  preservation  of,  320 

Wood  alcohol,  228 

Wood's  metal,  259 

Wool  fiber,  226 

Xenon,  84 
Yeast,  229 

Zinc,  317;   alloys  of,  360;  granu- 
lated, 318 ;  mossy,  318 
Zinc  chloride,  320 
Zinc  oxide,  319 
Zinc  sulfate,  320 
Zinc  white,  319 
Zincite,  317 


ANNOUNCEMENTS 


ELEMENTS  OF  GENERAL 
SCIENCE 

By  OTIS  W.  CALDWELL,  Head  of  the  Department  of  Natural  Science,  and 

WILLIAM  LEWIS  EIKENBERRY,  Instructor  in  the  University  High  School, 

School  of  Education,  The  University  of  Chicago 


8vo,  cloth,  xiv  +  308  pages,  illustrated,  $1.00 


THIS  book  presents  a  logically  arranged  and  teachable  first-year 
high-school  course  in  general  science,  dealing  with  concrete  scientific 
facts  of  everyday  interest  and  worth-while  significance.  The  material 
is  drawn  from  the  home,  school,  and  community  environment,  and 
all  the  sciences  contribute  to  the  survey,  each  being  used  as  needed. 
Unity  is  secured  through  the  logical  and  progressive  arrangement  of 
the  topics  which  make  up  the  course. 

These  fall  under  five  main  heads:  I.  The  Air,  II.  Water  and 
its  Uses,  III.  Work  and  Energy,  IV.  The  Earth's  Crust,  V.  Life 
upon  the  Earth.  Each  topic  is  connected  with  that  which  follows, 
the  last  in  one  main  division  leading  naturally  to  the  first  in  the  next 
division.  The  material  has  been  chosen  for  its  value  in  developing  a 
scientific  method  of  thinking,  and  in  giving  the  young  pupil  a  much- 
needed  basis  for  later  science  work  in  high  school.  Laboratory  work 
is  intended  to  accompany  the  text,  which  suggests  many  simple  but 
valuable  experiments. 

The  book  is  the  outcome  of  six  years'  experiment  with  general 
science  in  the  University  High  School,  The  University  of  Chicago. 
For  the  past  four  years  the  course  has  been  given  essentially  as  now 
published.  Its  plan  and  materials  have  in  addition  been  submitted  to 
many  high-school  teachers,  and  a  variety  of  helpful  suggestions  thus 
secured.  In  short,  the  book  has  been  made  in  the  laboratory  and  the 
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R?  A                             °-5 
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SEP     7  1940 

induction     1.60 

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UNIVERSITY  OF  CALIFORNIA  LIBRARY 


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II  II  II 


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£  J$ 
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LIST  OF  THE  ELEMENTS,  THEIR  SYMBOLS, 

AND  ATOMIC  WEIGHTS 

v 

The  more  important  elements  are  printed  in  heavier  type 


Aluminium 
Antimony  .     . 

Ar^on 

.     Al 
.     Sb 
A 

27.1 
120.2 
39.88 

Molybdenum    . 
Neodymium     . 
Neon 

.     Mo 
.     Nd 

Ne 

96.0 
144.3 

20.2 

Arsenic  . 

.     As 

74.96 

Nickel      .     .     . 

.     Ni 

58.68 

Barium 

.     Ba 

137.37 

Niton       .     .     . 

.     Nt 

222.4 

Bismuth      .     . 
Boron     .     .     . 

.     Bi 
.     B 

208.0 
11.0 

Nitrogen  .     .     . 
Osmium  .     .     . 

.     N 
.     Os 

14.01 
190.9 

Bromine      .     . 
Cadmium  . 

.     Br 

.     Cd 

79.92 
112.4 

Oxygen    .     .     . 

Palladium    . 

.     0 
.     Pd 

16.00 
106.7 

Caesium      .     . 
Calcium      .     . 

.     Cs 
.     Ca 

132.81 
40.07 

Phosphorus  .     . 
Platinum 

.     P 
.     Pt 

31.04 
195.2 

Carbon  . 

.     C 

12.00 

Potassium    . 

.     K 

39.10 

Cerium  . 
Chlorine      .     . 

.     Ce 

.     Cl 

140.25 
35.46 

Praseodymium 
Radium    . 

.     Pr 
.     Ra 

140.56 
226.4 

Chromium  . 

.     Cr 

52.0 

Rhodium      .     . 

.     Rh 

102.9 

Cobalt  .     .     . 

.     Co 

58.97 

Rubidium    . 

.     Rb 

85.45 

Columbium    . 

.     Cb 

93.5 

Ruthenium  . 

.     Ru 

101.7 

Copper  .     .     . 
Dysprosium    . 
Erbium" 

.     Cu 
.     Er 

63.57 
162.5 
167.7 

Samarium    .     . 
Scandium     . 
Selenium     .     . 

.     Sa 

.     Sc 
.     Se 

150.4 
44.1 
79.2 

Europium  .     . 
Fluorine 

Gadolinium    . 

.     Eu 
.     F 
.     Gd 

152.0 
19.0 
157.3 

Silicon     .     .     . 
Silver      .     .     . 
Sodium    .     .     . 

.     Si 
•     Ag 

.     Na 

28.3 
107.88 
23.00 

Gallium 

.     Ga 

69.9 

Strontium    . 

.     Sr 

87.63 

Germanium    . 

.     Ge 

72.5 

Sulfur      .     .     . 

.     S 

32.07 

Glucinum  . 
Gold  .... 

.     Gl 
Au 

9.1 
197.2 

Tantalum     . 
Tellurium    . 

.     Ta 
.     Te 

181.5 
127.5 

Helium      .     . 

.     He 

3.99 

Terbium       .     . 

.     Tb 

159.2 

Holmium   . 

.     Ho 

163.5 

Thallium     .     . 

.     Tl 

204.0 

Hydrogen    .     . 
Indium 
Iodine 

.     H 
.     In 
.     I 

1.008 
114.8 
126.92 

Thorium      .     . 
Thulium       .     . 
Tin     . 

.     Th 
.     Tm 

Sn 

232.4 

168.5 
119.0 

Iridiurri 
Iron  . 

.     Ir 
Fe 

193.1 
55.84 

Titanium 
Tungsten 

.     Ti 
.     W 

48.1 
184.0 

Krypton     .     . 
Lanthanum    . 

.     Kr 
.     La 

82.92 
139.0 

Uranium      .     . 
Vanadium 

.    u 

.     V 

238.5 
51.0 

Lead       .     .     . 

.     Pb 

207.1 

Xenon     .     .     . 

.     Xe 

130.2 

Lithium 

.     Li 

6.94 

Ytterbium    . 

.     Yb 

172.0 

Lutecium  .     . 

.     Lu 

174.0 

Yttrium  .     .     . 

.     Yt 

89.0 

Magnesium 
Manganese 
Mercury      .     . 

.     Mg 
.     Mn 
.     Hg 

24.32 
54.9:5 
200.6 

Zinc    .... 
Zirconium    .     . 

.  *  -i    O</ 

.     Zn 
.     Zr 

65.37 
90.6 

